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REESE   LIBRARY 


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UNIVERSITY  OF  CALIFORNIA. 
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EXPERIMENTS 


ARRANGED    FOR 


STUDENTS  IN  GENERAL  CHEMISTRY 


BY 


EDGAR  F.  SMITH       AND        HARRY  F.  KELLER, 
it 

PROFESSOR     OF     CHEMISTRY,     UNIVERSITY     OF  PROFESSOR    OF    CHEMISTRY,    MICHIGAN   MINING 

PENNSYLVANIA,    PHILADELPHIA.  SCHOOL,  HOUGHTON. 


SECOND  EDITION,  ENLARGED,  WITH 37  ILLUSTRATIONS. 


PHILADELPHIA: 

P.    BLAKISTON,    SON    &    CO 

1012    WALNUT     STREET. 


|)Q$ 


COPYRIGHT,  1891,  BY  P.  BLAKISTON,  SON  &  Co. 


PRESS  or  WM   F.  FELL  &  Co., 
I22O-24  SANSOM  ST., 

PHILADELPHIA. 


PREFACE. 


This  little  work  is  designed  as  a  guide  for  beginners  in  chemistry. 
The  arrangement  of  the  course  is  such  as  the  authors  have  used  with 
success  in  the  instruction  of  their  classes ;  its  object  is  not  to  dispense 
with  the  supervision  of  an  instructor,  but  rather  to  assist  him.  Although 
reference  is  made  to  Richter's  "Inorganic  Chemistry,"  any  other  text- 
book on  the  subject  can  be  employed  in  its  stead.  The  experiments 
have  been  collected  from  various  sources,  and  no  claim  is  made  for 
originality. 


CONTENTS. 


PART  I. 

CHAPTER  PAGE 

I.  APPARATUS,  MANIPULATIONS  AND  OPERATIONS, 9-10 

II.  GENERAL  PRINCIPLES     10-11 

III.  HYDROGEN, 12-14 

IV.  CHLORINE,  BROMINE,  IODINE,  FLUORINE, 14-19 

V.  OXYGEN,  SULPHUR, .— .          19-25 

VI.  NITROGEN,  PHOSPHORUS,  ARSENIC,  ANTIMONY,      ." 25-32 

VII.  CARBON  AND  SILICON,  BORON, 32-34 

PART  II. 

VIII.  POTASSIUM,  SODIUM  [AMMONIUM], 35-38 

IX.  CALCIUM,  STRONTIUM,  BARIUM, 39-41 

X.  MAGNESIUM,  ZINC, 41-42 

XL  MERCURY,  COPPER,  SILVER,  GOLD, 43-46 

XII.  ALUMINIUM,  TIN,  LEAD,  BISMUTH,      46-50 

XIII.  CHROMIUM,  MANGANESE,  IRON,  NICKEL,  COBALT, 50-56 


NON-METALS. 


FIG.  i. 


CHAPTER  I. 

APPARATUS,  MANIPULATIONS  AND  OPERATIONS. 

1 i )  The  Bunsen  burner  and  the  blowpipe. 

i.  Make  a  borax  bead.  2.  Dissolve  a  very  minute  quantity  of  man- 
ganese dioxide  in  it.  3.  Heat  in  the  oxidizing  flame  (?).  4.  In  the 
reducing  flame  (?).  5.  Heat  oxide  of  lead  on  charcoal  in  the  reducing 
flame.  6.  In  the  oxidizing  flame. 

(2)  Working  with  glass  tubing  and  rods. 

i.  Cut  various  lengths  of  rods  and  tubing.  2.  Round  the  sharp  edges 
by  softening  and  turning  the  ends  in  the  lamp. 

(3)  Construct  a  wash-bottle  (Fig.  i). 

i.  Soften  a  sound  cork  by  rolling  it  under  your  foot  on 
a  clean  floor.  2.  Bore  two  parallel  holes  through  it  by 
means  of  a  cork-borer.  These  perforations  should  be  cyl- 
indrical and  of  less  diameter  than  the  glass  tubes  they  are 
to  receive.  Use  a  rat-tail  file  in  enlarging  them.  3.  Cut 
suitable  lengths  of  glass  tubing.  4.  Draw  the  longer  one 
to  a  fine  point  after  softening  in  the  flame.  5.  Bend  the 
tubes  in  an  ordinary  fish-tail  burner,  and  round  the  sharp 
edges.  6.  Fit  the  different  pieces  together. 

(4)  Arrange  some  other  form  of  apparatus  for  practice. 

(5)  The  balance. 

i.  Weigh  an  object  by  placing  it  on  the  left-hand  pan  of  the  balance, 
and  a  weight  judged  about  equal  on  the  right-hand  pan.  Should  the 
latter  be  found  too  heavy,  replace  it  by  the  next  smaller  one  \  if  too  light, 
by  the  next  heavier  one.  Then  add  systematically  the  smaller  weights, 
until  the  needle  points  to  the  middle  of  the  scale.  The  final  adjustment 
is  made  with  the  rider.  In  adding  or  removing  weights,  the  supports 
must  always  be  raised. 

(6)  Measuring  vessels. 

i.  Measure  off  10  cc.  of  water  (a]  in  a  cylinder,  (£)  in  a  burette,  (c) 

9 


10  EXPERIMENTS    IN    GENERAL    CHEMISTRY. 

in  a  pipette.  Always  read  the  lower  meniscus.  2.  Measure  off  similarly 
50,  100  and  200  cc.  of  water,  and  determine  their  weight.  3.  Measure 
the  volume  of  50  grms.  of  oil  of  vitriol,  and  of  65  grms.  of  muriatic  acid. 
What  are  the  specific  gravities  of  these  substances  ?  Note  the  relation 
between  weight  and  volume  in  the  metric  system. 

(7)  Chemical  operations  :  Solution,  evaporation,  crystallization,  precip- 
itation, filtration,  washing  and  drying. 

i.  Place  into  a  test-tube  pure  sodium  carbonate,  into  another  cobalt 
chloride,  and  add  distilled  water  to  each.  Stir.  What  occurs  ?  2.  To 
calcium  carbonate,  add  water.  Is  there  any  change  ?  Now  add  a  little 
hydrochloric  acid.  What  action  has  it  ?  3.  Pour  5  cc.  of  strong  hydro- 
chloric acid  upon  powdered  manganese  dioxide ;  observe  appearance  and 
odor.  Note,  too,  in  each  case,  whether  heat  has  any  effect.  Distinguish 
between  chemical  and  mechanical  solution.  4.  Heat  the  cobalt  chloride 
and  the  calcium  carbonate  solutions,  each  in  a  separate  dish,  on  an  iron 
plate,  until  the  liquids  are  completely  driven  off  (?).  5.  Dissolve  potas- 
sium chlorate  in  hot  water,  and  allow  to  stand  and  cool  (?).  6.  To  a 
portion  of  the  cobalt  chloride  solution,  add  a  solution  of  soda;  boil.  7. 
Allow  to  settle  and  filter.  8.  Wash  the  precipitate  "until  pure  water  runs 
through  the  filter  (?).  9.  Heat  the  filter  until  perfectly  dry. 


CHAPTER  II. 

GENERAL    PRINCIPLES. 

(i)   Changes  in  matter. 

i.  Rub  a  glass  rod  with  a  piece  of  cloth,  then  touch  particles  of  paper 
with  it  (?).  2.  Through  an  insulated  spiral  of  stout  copper  wire  pass  a 
current  from  two  Bunsen  cells.  Place  a  piece  of  wrought-iron — a  nail 
will  answer — inside  the  spiral,  and  bring  iron  filings  in  contact  with  it. 
What  happens?  Interrupt  the  current  and  note  the  result ;  repeat.  3. 
Heat  a  platinum  wire  in  the  non-luminous  flame  ;  is  there  any  change  ? 
What  is  the  effect  of  removing  it  ? 

Are  the  original  properties  of  the  substances  in  the  above  experiments 
altered,  after  the  action  of  the  forces  of  electricity,  magnetism  and  heat 
has  been  stopped  ? 

4.  Mix  intimately  four  parts,  by  weight,  of  finely  powdered  sulphur 
with  seven  parts  of  very  finely  divided  iron  (filings).  Pass  a  magnet  over 
a  portion  of  the  mixture.  Another  portion  treat  with  carbon  disulphide  in 


GENERAL   PRINCIPLES.  II 

a  test-tube.  Then  heat  the  remaining  portion  in  a  tube  over  a  gas  flame. 
Note  carefully  what  occurs  in  each  case.  Powder  the  mass  FIG  a 
resulting  from  the  last  operation  in  a  dry  mortar.  Can  you 
extract  from  it  any  iron  with  a  magnet,  or  dissolve  out  any 
sulphur  with  carbon  disulphide?  What  inference  do  you  draw 
from  the  facts  observed  ?  5.  Decompose  water  in  Hofmann's 
apparatus  by  an  electric  current.  The  water  should  be  acid- 
ulated with  sulphuric  acid  to  make  it  a  conductor  of  elec- 
tricity. A  current  from  four  to  six  Bunsen  cells  is  required. 
To  the  gas,  of  which  a  larger  volume  has  collected,  apply  a  flame,  and  to 
the  other  a  glowing  spark  (?)  6.  Heat  oxide  of  mercury  in  a  tube  of 
hard  glass  (Fig.  2).  Apply  the  spark  test  (?).  7.  Rub  some  sulphur 
and  mercury  together  in  a  mortar  (?)  8.  Heat  sugar  in  a  dry  test-tube, 
at  first  gently,  and  then  more  strongly.  Note  color  and  odor.  9.  Mix 
dry  soda  and  tartaric  acid  in  a  mortar.  Is  there  any  action  ?  What 
occurs  when  you  add  water  ?  Point  out  in  what  respect  the  changes 
involved  in  experiments  1-3  differ  essentially  from  those  in  4-9.  By 
what  general  names  can  you  distinguish  the  two  different  kinds?  With 
which  does  chemisty  concern  itself?  Define  chemistry. 

Through  what  agencies  have  the  results  been  obtained  in  experiments 
4-9  ?     Has  any  gain  or  loss  of  matter  occurred  in  any  of  them  ? 

(2)  The  products  resulting  from  5  and  6  cannot  be  further  simplified, 
*.  e.,  decomposed  into  dissimilar  substances.     They  are  elements*     What 
are  water  and  red  oxide  of  mercury  ? 

i.   Dissolve  in  a  little  nitric  acid,  the  black  powder  obtained  by  heating 
an  intimate  mixture  of  powdered  sulphur  and  copper  filings. f 

Evaporate  the  solution  nearly  to  dryness,  take  up  in  water       FIG.  3. 
and  filter.     What  remains  on  the  filter  ?     Place  the  filtrate  in 
a  beaker,  dip  the  platinum  electrodes  of  a  battery  (one  or  two    ~*"P 
Bunsen  cells)  into  it  (Fig.  3),  and  allow  the  current  to  act  for 
ten  minutes.     What  do  you  observe  upon  the  platinum  foil, 
forming  the  negative  pole  ?     What  changes  have  the  copper 
and  the  sulphur  undergone  in  this  experiment  ? 

(Study  pp.  18-27,  in  Richter's  Chemistry.) 

(3)  Metals  and  non-metals.     (See  Richter,  p.  20.) 


*  The  instructor  should  here  develop  the  idea  of  element  more  fully. 

f  A  better  substitute  would  be  finely  divided  copper ;  such  as  may  be  obtained  by  the 
reduction  of  black  cupric  oxide  in  a  current  of  hydrogen  gas  (see  page  14). 


12 


EXPERIMENTS    IN    GENERAL   CHEMISTRY. 


CHAPTER   III. 

HYDROGEN.— H. 

(i)  Put  several  pieces  of  granulated  zinc  into  a  test-tube  and   pour 
dilute  sulphuric  acid  upon  them.     What  occurs? 

FlG.  4.  (2)  Arrange    the   apparatus   shown   in   Fig.   4. 

The  flask  should  contain  about  15  grms.  of  Zn, 
and  dilute  H2SO4  is  poured  through  the  funnel 
tube.  When  all  the  air  in  the  apparatus  has  been 
displaced  (ask  for  precautions  /)  collect  six  test- 
tubes  full  of  the  gas  over  water. 

(3)  What  are  its  properties  ?  Will  it  burn  ? 
Support  combustion?  Is  it  lighter  than  air? 
(4)  i.  To  learn  what  becomes  of  hydrogen  when  it  burns  in  air, 
arrange  apparatus  as  in  Fig.  5.  The  gas  is  led  from  the  evolution  flask 
A,  into  a  bottle  containing  concentrated  H2SO4,  and  then  passes  through 
a  tube  filled  with  pieces  of  CaCl2.  The  gas  which  escapes  is  free  from 
moisture.  Burn  it  under  a  cold  glass  jar.  What  do  you  obtain  ?  2.  Fill  a 
small  flask  with  a  mixture  of  one  vol.  of  H  and  five  vols.  of  air ;  cork ; 
invert  the  flask  several  times  to  mix  the  gases;  wrap  a  towel  around  it  and 
bring  its  mouth  to  a  flame.  Result  ? 

FIG.  5.  (5)  Hydrogen  is  not  the  only  product  of 

the  action  of  Jf2SO4  itpo?i  Zn. 

Pour  some  of  the  liquid  remaining  in 
the  flask,  in  which  H  was  generated,  into 
a  porcelain  dish.  Evaporate  to  about 
one-third  of  the  original  bulk ;  allow  to 
stand  several  hours.  You  will  now  dis- 
cover that  the  solution  is  full  of  colorless 
crystals.  These  are  zinc  sulphate  or  white  vitriol — a  salt,  ZnSO4  -j-  7H2O. 
Write  the  equation  of  the  reaction. 

(6)  Determine  the  weight  of  H  generated  by  a  given  weight  of  Zn. 

A  piece  of  Zn  (not  more  than  .02  grm.)  is  accurately 
weighed,  and  placed  under  a  funnel  in  a  beaker  (Fig.  6). 
The  latter  is  then  nearly  filled  with  water,  so  that  the  en- 
tire funnel  is  under  the  surface.  A  test-tube  containing  dilute 
H,SO4  is  lowered  over  the  stem  of  the  funnel.  Hydrogen 
appears  and  collects  in  the  tube.  When  all  the  Zn  has  dis- 
appeared,* transfer  the  tube  containing  the  H  to  a  larger  vessel, 
holding  water.  Measure  the  volume  of  the  gas  by  marking 


FIG.  6. 


This  may  be  hastened  by  bringing  a  spiral  of  platinum  wire  in  contact  with  the  Zn. 


HYDROGEN.  1 3 

the  tube  where  the  inner  and  outer  levels  of  water  are  even,  and  then 
weighing  or  measuring  the  quantity  of  water  that  it  will  hold  to  that 
mark.  Note  the  temperature  of  the  water,  and  the  height  of  the  baro- 
meter. 

The  weight  of  the  H  is  found  by  multiplying  the  vol.  by  the  wt.  of 
i  cc.,  /.<?.,  .0000896  gr.  Before  this  can  be  done,  however,  it  is  necessary 
to  reduce  the  volume  of  the  gas  to  o°  C.  and  760  mm.,  as  the  above  value 
has  been  determined  under  these  conditions.  If  v  =r  volume  observed, 
t  —  temperature,  and  p  =  pressure,  then 


v  X  p 


(i  +  at)  X 


and  W 


X  .000x5896  . 


a  =  .003665.* 


To  calculate  the  quantity  of  Zn  necessary  to  generate  a 
unit  of  H,  we  say — 


FIG.  7. 


Wt  of  H  :  Wt  of  Zn  :  :  I  :  x. 

x  here  stands  for  the  equivalent  weight  of  Zn. 

The  equivalent  weights  of  some  other  metals,  such  as 
Fe,  Cd,  Mg,  can  be  determined  in  the  same  manner.  Mag- 
nesium gives  the  most  satisfactory  results. 

(7)  Decompose  water  by  electrolysis  and  test  the  products. 

(8)  Take  a  small  piece  of  sodium,  wrap  it  in  paper  and  place  it,  with 
forceps,  under  the  mouth  of  a  test-tube 

filled  with  water,  and  inverted  in  water 
(Fig.  7)  contained  in  a  dish.  Repeat  this 
until  the  test-tube  is  filled  with  the  gas. 
Test  it  for  H. 

What  becomes  of  the  metal  ?     Write    n 
the  reaction. 

(9)  Construct  the  apparatus  shown  in 
Fig.  8. 

Water  is  heated  to  boiling  in  the  flask 

A,  and  the  steam  led  over  iron  filings  or  wire,  heated  to  redness  in  a  tube 
of  hard  glass.  Care  must  be  taken  to  prevent  steam  from  condensing  in 
any  part  of  the  tube.  Collect  the  escaping  gas  over  water.  Test  it 
for  H. 


*  Tension  of  aqueous  vapor  is  here  neglected. 


EXPERIMENTS    IN    GENERAL   CHEMISTRY. 


Is  the  iron  changed  ?     Equation  ? 

FIG  (10)  Into    a   weighed    tube 

of  hard  glass  (6-8  inches  in 
length)  place  a  weighed  quan- 
tity (1-2  grams)  of  cupric 
oxide ;  connect  the  tube  with 
a  CaCl2-tube  of  known  weight 
(Fig.  9).  Pass  a  current  of  dry 
H  over  the  CuO  and  heat. 
After  the  change  is  complete,  cool  and  determine  the  loss  in  weight  of 
tube  -)-  CuO,  and  the  gain  in  the  CaCL2  tube.  Explain  the  reaction. 

Problems. — i.   How  much  H  can  be  obtained  from  Zn  and   299  grms. 
of  H2SO4?     2.   How  much   Zn  and  H-jSO^  are  necessary  to  furnish   100 
grms.  of  H  ?     3.   Suppose   you  have  found  that  .015  grm.  of  Mg  yield 
15.2  cc.  of  H  at  20°  C.  and    750  mm.,  what  is  the  equivalent  weight  of 
that  metal  ?     4.   How  many  cc.    of  H  can  be  obtained   from  2  grms.  of 
Na  and  water?     5.  How  many  grms.   of  H2O  can  be  decomposed  by 
5  grms.  of  Fe;  by  how  much  is  the  weight  of  the  latter  increased  ?     6. 
10  grms.  of  CuO  will  yield  how  much  Cu  upon  heating  in  H? 
Give  a  brief  summary  of  what  you  have  learned  about  H. 
(Study  Richter,  pp.  39-47-) 


FIG. 


CHAPTER   IV. 

FIRST  NATURAL  GROUP  OF  ELEMENTS— CHLORINE,  BROMINE,  IODINE, 

FLUORINE. 

CHLORINE.— Cl. 

(i)  Into  a  test-tube  put  MnO2  and  concentrated  HC1.     Note  what 
happens  both  before  and  after  heating. 

(2)  Use  apparatus  (shown  in  Fig.  10)  for  preparing 
larger  quantities  of  Cl.     The  MnO2  should  be  in  the 
form  of  small  lumps  (not  powder.)     Heat  the  mix- 
ture gently,  pass  the  Cl  through  a  small  quantity  of 
water  and  collect  it  either   by  downward    displace- 
ment or  over  warm  water. 

Write  the  reaction.     How  many  atoms  of  Cl  are 
liberated  ?     How  many  molecules  ? 

(3)  i.  What  is  the  normal  condition  of  this  ele- 
ment?    2.   Is  it  lighter  than  air?     3.   Is  it  inflamma- 

4.   Does  it  support  combustion? 


ble? 


FIRST    NATURAL   GROUP   OF    ELEMENTS — CHLORINE.  15 

To  obtain  answers  to  these  questions,  fill  5  test-tubes  with  dry  Cl,  and 
proceed  as  under  H. 

(4)  Again  fill  5  bottles  with  the  dry  gas.     Cover  them  with  glass  plates. 
Into  i  throw  a  little  pulverized  antimony. 

Into  2  carefully  introduce  a  piece  of  phosphorus. 

Into  3  insert  tissue  paper  moistened  with  oil  of  turpentine.  * 

Into  4  introduce  colored  flowers. 

Into  5  pour  a  little  litmus  solution. 

What  are  the  results  ? 

(5)  Fill  a  small-sized  flask  one-half  with  chlorine,  the  other  half  with 
hydrogen.     Wrap  a  towel  about  the  flask  and  apply  a  flame  to  its  open 
mouth  (?).      Care. 

(6)  Invert  a  bottle  filled  with  Cl  over  water  saturated  with  the  same 
gas.     What  follows  in  the  course  of  a  few  hours'  exposure  to  sunlight  ? 
Can  you  account  for  results  in  experiments  (4),  (5)  and  (6)  ?     Why  should 
the  gas  be  collected  over  warm  water  ? 

(7)  Determine  the  weight  of  a  litre  of  chlorine.     Arrange    apparatus 
as  shown  in  Fig.  n. 

In  the  evolution-flask  place  a  mixture  of  equal 
weights  of  salt  and  manganese  dioxide.  Add  sul- 
phuric acid,  previously  diluted  with  its  own  volume 
of  water  (pour  the  acid  into  the  water !  ).  Heat 
gently.  Chlorine  is  evolved,  and  dried  by  passing 
it  through  concentrated  sulphuric  acid,  after  which  it 
is  led  into  the  perfectly  dry  flask  c.  When  this  is 
filled,  which  you  ascertain  by  the  color  of  the  gas  in 
the  neck,  slowly  withdraw  the  tube  and  cork  the 
flask  at  once.  Weigh  the  flask.  Read  the  barometer  and  thermo- 
meter. Determine,  also,  the  weights  of  the  flask  filled  with  air  and  with 
water. 

Calculation : 

Capacity  of  flask, a 

Temperature, t 

Pressure,     p 

Flask  filled  with  air, .    w 

Cl, \    .  w' 

Wt.  of  a  litre  of  air, *-293  grm. 

"        Cl,     ....'.'.....    x 

*  It  is  well,  when  the  turpentine  is  old,  to  gently  warm  it,  and  then  saturate  the  tissue 
paper. 


1 6  EXPERIMENTS    IN    GENERAL   CHEMISTRY. 

The  weight  of  the  air  filling  the  flask  is    axPx-OOI293  ^     The 

(i  +  .00367  t)  760 

ence  between  this  and  w  is  the  weight  of  the  vacuous  flask.     Subtract 
this  from  w'.  The  remainder  is  the  weight  of  the  Cl,  (W).   Reduce  the  vol. 

of  the  Cl  to  o°  C.  and  760  mm.  (see  under  H) ;  it  is  v0  = ax  p 

(  i  +  .003671)760 

11  •     i  ^       r        T^  W  X    IOOO 

and  the  weight  of  i  litre,  x  =  -    — 

How  much  heavier  is  one  litre  of  Cl  than  an  equal  vol.  of  H  ? 

Write  the  reaction  involved  in  the  above  method  for  preparing  Cl. 

Problems. — i.  How  many  litres  of  Cl  can  be  obtained  from  i  kilo  of 
MnO2  and  HC1?  2.  What  weight  of  salt  is  required  to  prepare  100 
litres  of  Cl  ?  3.  How  many  pounds  of  sodium  sulphate  and  manganese 
sulphate  will  be  formed  in  the  preparation  of  100  litres  of  chlorine  gas? 
4.  Calculate  the  number  of  grams  of  Cl  that  2  litres  of  water  will  ab- 
sorb, provided  the  latter  takes  up  twice  its  volume  of  Cl  ?  Write  out 
your  deductions  from  the  above  experiments  on  Cl. 

(Read  Richter,  pp.  49-52.) 


HYDROGEN  CHLORIDE.— HC1. 

(1)  Repeat  the  explosion  of  equal  vols.  of  Cl  and  H.     Quickly  cover 
the  mouth  of  the  flask,  and  immerse  it  under  water.     Does  the  latter  rise  ? 
Put  a  drop  of  the  liquid  on  the  tongue  and  note  the  taste.     Add  some 
blue  litmus  solution.     Is  there  any  change? 

(2)  The  product  of  the  union  of  H  and  Cl  is  a  colorless  gas.     It  is 
called  hydrogen  chloride.     It  is  usually  prepared  by  the  action  of  sulphuric 
acid  upon  salt,  thus: — 

2NaCl  +  H2S04  =  2HC1  +  Na2SO4. 
or,  better,  NaCl  +  H2SO4  =  HC1  -f  NaHSO4. 

The  apparatus  employed  here  is  the  same  as  that  used  for  making  Cl 
(Fig.  10). 

(3)  Determine  the  properties  of  HCl  as  under  H  and  CL 

What  new  property  appears  here  ?  Fill  a  long  dry  glass  tube  with  the 
gas,  and  quickly  bring  it  into  a  basin  containing  water  colored  blue  with 
litmus.  What  happens?  What  does  HCl  gas  yield  on  dissolving  in 
water  ? 

(4)  In  the  preparation  of  H  by  the  action  of  Na  upon  water,  it  was 
observed  that  the  liquid  became  soapy  to   the   touch,  and  acquired  the 
property  of  turning  red  litmus  blue.     Prepare  such  a  solution.     To  it  add 
a  few  drops  of  litmus,  and   then   an   HCl  solution   (gradually)   from   a 


FIG.  12. 


FIG.  13. 


FIRST    NATURAL    GROUP   OF    ELEMENTS — HYDROGEN    CHLORIDE.          17 

burette,  until  the  blue  color  just  begins  to  turn.     Evaporate  the  resulting 

liquid   to  crystallization.      Dissolve  and   recrystallize  the  product.      It 

appears  in  cubes,  and  has  the  taste  of  common  salt.     It 

does  not  affect  either  red  or  blue  litmus.     We  say  it  is 

neutral  in  reaction.     The  substance  is  chloride  of  sodium 

or  common  salt.     What  is  a  salt?     An  acid?     A  base  ? 

How  can  you  obtain  HC1  and  Cl  from  NaCl  ? 

(5)  Burn Hin  an  atmosphere  of  Cl,  and  Clin  hydrogen. 
Generate  chlorine  as  already  described  (p.  14)  and 

collect  it  in  a  large  cylinder.     Into  this  introduce  a  burn- 
ing jet  of  hydrogen  (Fig.  12).     Does  it  continue  burn- 
ing ?     What  is  the  appearance  of  the  flame  ?     To  show  the  combustion  of 
Cl  in  H  arrange  apparatus  as  in  Fig.  13. 

(6)  To  determine  the  weight  of  a  litre   of  HCl,   proceed 
exactly  as  under  chlorine. 

(7)  Determine  the  composition  of  HCl  by  volume. 

i.  Fill  a  perfectly  dry  and  graduated  tube  with  HCL  Close 
the  open  end  with  the  thumb,  and  opening  the  tube  for  a 
moment,  quickly  pour  in  about  10  cc.  of  sodium  amalgam 
(see  sodium,  p.  36).  Close  the  tube  at  once  with  the  thumb, 
slightly  moist,  and  shake  well.  Invert  the  tube  in  a  large 
beaker  of  water,  and  remove  the  thumb.  The  amalgam  will 
drop  into  the  water,  and  the  latter  will  rush  up  into  the  tube, 
filling  it  nearly  half  full.  Immerse  the  tube  so  that  the  water 
in  it  and  that  in  the  beaker  are  on  the  same  level.  This  is 
done  to  measure  the  hydrogen  under  atmospheric  pressure. 
Read  the  residual  volume  of  the  gas  and  measure  the  volume 
of  the  mercury. 

Calculation : — 

Capacity  of  tube, a 

Vol.  of  mercury,       b 

Vol.  of  H c 

a  — b 


(8)  Add  a  solution  of  HCl  to  solutions  of  silver  nitrate  ;  of  mercurous 
nitrate  ;  and  of  lead  acetate.  What  do  you  observe  in  each  case?  Boil 
the  precipitate  formed  in  the  lead  solution  with  water.  Cool,  and  note 
result. 

3 


i8 


EXPERIMENTS    IN    GENERAL    CHEMISTRY. 


FIG.  14. 


BROMINE.— Br. 

(1)  Allow  a  drop  of  bromine  to  fall  upon  a  heated  watch  glass  ;  cover 
it  quickly  with  a  beaker.     What  is  the  color  of  the  vapor  ?     Dissolve  one 
drop  of  bromine  in  each  of  the  following  solvents  contained  in  test-tubes : 
water,   alcohol,   ether,  carbon   disulphide,   and  chloroform.      Note    the 
relative  solubilities,  and  the  color  of  each  solution. 

(2)  i.   Pass  Cl   through  an   aqueous  solution  of   potassium  bromide. 
What  happens?     2.   To  one  portion  of  the  product  add  a  few  drops  of 

CS2,  and  agitate  the  mixture ;  what  is  the  result  ? 
3.  To  another  portion  of  the  solution,  containing 
free  Br,  add  a  few  drops  of  a  starch  solution.* 
Result  ? 

(4)  Devise    a   method    for   preparing  bromine 
from  KBr. 

(5)  Prepare  hydrobromic  acid. — In  a  small  flask 
cover  2  grams  of  amorphous    phosphorus  with  4 
grams  of  H2O,  and  from  a  funnel,  provided  with 

a  stop-cock,  gradually  allow   20  grams  of  Br  to   run-  in.f 

The  gas  is  purified  by  conducting  it  through  a  U-tube,  containing  moist- 
ened pieces  of  phosphorus  and  glass  (Fig.  14),  and  led  into  water  to 
obtain  the  aqueous  solution. 

(6)  Add  aqueous  HBr  to  solutions  of  AgNO3,  HgNO3  and  Pb(NO3)2— 
do  the  resulting  bromide  precipitates  differ  much  from  the  corresponding 
chlorides  ? 

IODINE.— I. 

(1)  i.  Place  an  iodine  crystal  upon  a  warm  plate,  and  note  color  of 
vapor.     2.  Test  the  solubility  of  iodine  in  the  same  solvents  as  were  used 
with  bromine  ;  what  are  the  colors  of  the  resulting  solutions? 

(2)  i.   Pass  Cl  through  a  solution  of  KI.     Test  the  resulting  liquid 
with  ether,  carbon  disulphide  and  starch  solution  (as  with  Br).      2.  Repeat 
this  experiment,  substituting  Br- water  for  the  Cl.     Avoid  excess  of  Cl  as 
well  as  Br  (?). 

What  conclusion  do  you  draw  from  these  experiments  relative  to  the 
affinity  of  the  halogens  for  potassium  ? 

*  The  starch  solution  necessary  for  this  purpose  can  be  prepared  as  follows  :  One  gram 
of  starch  is  well  ground  in  a  mortar,  with  very  little  water,  to  creamy  consistence.  It  is 
then  poured  into  200  cc.  of  boiling  water.  Allow  to  subside,  decant  the  clear  superna- 
tant liquid  and  use  it  for  the  test. 

f  As  it  is  rather  difficult  to  weigh  bromine  upon  a  balance,  calculate  the  volume  corres- 
ponding to  the  weight  given  and  measure  out  the  same  in  a  cylinder. 


SECOND  NATURAL  GROUP  OF  ELEMENTS — OXYGEN.          19 

(3)  Pass  hydrogen  sulphide  gas  (H2S)   into   50  cc.  of  water,  and  add 
powdered  iodine  till  the  brown  color  no  longer  disappears.     Warm,  filter 
(?)  and  distil  the  filtrate.     The  product  is  what  ? 

How  is  gaseous  HI  prepared  ? 

(4)  Precipitate  solutions  of  silver  nitrate  (AgNO3),  mercurous  nitrate 
(HgNO3),  lead  nitrate  (Pb(NO3)2),  and  mercuric  chloride  (HgCl2)  with 
KI.     Note  result  in  each  case. 


FLUORINE.— Fl. 

(1)  In  a  lead  dish  (or  platinum  crucible)   place  i  gram  of  pulverized 
fluor  spar  (CaFl2).     Add  cone.  H2SO4 ;  cover  the  dish  or  crucible  with  a 
watch-glass  coated  with  paraffin,  through  which  some  characters  have  been 
drawn  with  a  fine  point.     Heat  gently  for  a  few  minutes. 

What  do  you  observe  on  removing  the  paraffin  ? 

(2)  Can  you  liberate  Fl  from  a  fluoride? 

Problems. — i.  How  much  NaBr,  H2SO4and  MnO2are  necessary  to  pro- 
duce i  cu.  metre  of  Br  vapor  at  20°  C  and  745  mm.  ?  2.  What  per  cent,  of 
HI  does  a  liquid  contain,  which  represents  a  solution  of  50  litres  of  the 
gas  in  i  litre  of  H2O?  3.  10  grms.  of  CaFl2  will  give  what  weight  of 
HF1  ?  4.  How  much  salt  and  sulphuric  acid  will  be  required  to  prepare 
6  litres  of  muriatic  acid  of  sp.  gr.  1.17?  What  volume  would  the  HC1 
in  these  six  litres  occupy  at  735  mm.  pressure  and  22°C?  5.  What  is 
the  percentage  of  hydrochloric  acid  in  a  solution  of  which  17  cc.  dis- 
solve exactly  2  grams  of  metallic  magnesium  ?  What  is  the  volume  of 
hydrogen  liberated  at  760  mm.  and  o°  ? 


CHAPTER  V. 

SECOND  NATURAL  GROUP  OF  ELEMENTS— OXYGEN,  SULPHUR, 
SELENIUM,  TELLURIUM. 

OXYGEN.— O. 

(1)  Preparation. — i.  Weigh  the  hard  glass  tube  a  (Fig.  15),  and  intro- 
duce a  weighed  quantity  (about  .5  grm.)  of  red  oxide  of  mercury.    Ignite 
strongly;    collect  the  liberated  gas,  and   measure   it.     Weigh    the  tube 
with  the  residue.     What  are  the  products  of  the  ignition  ? 

(2)  Prepare  more  of  the  gas,  as  follows :     Mix  equal  parts  of  KC1O3 
and   pulverized    MnO2 ;    heat  in  a  tube  of  hard  glass  or   small  retort. 
Collect  the  gas  in  bottles  over  water  (Fig.  15). 


20 


EXPERIMENTS    IN    GENERAL   CHEMISTRY. 


FIG.  15. 


Into  No.  i  lower  a  piece  of  ignited  sulphur  on  an  iron  spoon.     Note 
result.     Add  water  after  the  combustion  (?). 

Into  No.  2  introduce  a  small  piece  of  burn- 
ing phosphorus  (care  !).  Proceed  as  in  No.  i. 

Into  No.  3  introduce  ignited  charcoal.  Treat 
as  before.  Add  now  a  few  drops  of  blue  litmus 
to  the  contents  of  each  bottle.  Any  change? 

Into  bottle  No.  4  introduce  a  fine  watch 
spring,  previously  heated  at  one  end  and 
dipped  into  powdered  S.  Result  ? 

Is  oxygen  heavier  or  lighter  than  air? 
it  color,  taste,  or  odor  ? 
it  support  combustion  ? 


Has 

Will  it  burn  ?     Does 


What  other  methods  can  be  used  for  preparing  O  ? 
(3)  Determine  the  weight  of  a  litre  of  O. 


FIG. 


Arrange  apparatus  shown  in  Fig.  16;  a  is  a 
tube  of  hard  glass,  whose  weight  is  known ;  it 
contains  a  weighed  amount  of  KC1O3  (about  0.3 
grm.).  The  bottle  A  is  filled  with  water,  b  is  a 
clip  and  d  a  beaker.  The  exit  tube  should  con- 
tain water  as  far  as  clip  b  at  the  beginning  of  the 
experiment.  Open  the  clip,  heat  a  to  bright 
redness,  and  receive  the  water  displaced  by  the 
O  in  d.  When  no  more  gas  is  evolved,  cool ;  allowing  the  rubber  tube 
to  dip  under  the  water  of  the  jar.  Some  of  the  water  will  be  drawn  back 
into  the  bottle  (?).  Measure  the  volume  of  the  water  in  d.  Note  the 
temperature  of  the  air,  and  the  height  of  the  barometer.  Weigh  a,  con- 
taining residue  (KC1). 

Calculation  : 

Weight  of  the  tube, a 

Weight  of  KC1O3  and  tube, b 

Weight  of  KC1O3, b  —  a 

Volume  of  H2O  collected, v 

Barometric  pressure, p 

Temperature, t 

Aqueous  tension  at  t,     ...        p7 

Weight  of  KC1  and  tube c 

VX(P  —  P')  (b— c)xiooo 


Vo_. 


"(I +  .003670x760 


and 


.Vo 


Dissolve  the  residual  KC1  in  water,  and  to  its  solution  add  nitrate  of 
silver  (?).     How  does  KC1O3  behave  under  like  conditions? 


SECOND    NATURAL   GROUP   OF    ELEMENTS WATER.  21 

(4)  Give  a  summary  of"  your  work  upon  O. 

Problems. — i.  How  much  O,  by  wt.  and  vol.,  can  be  obtained  from 
54  grms.  of  HgO?  2.  Heat  will  expel  what  vol.  of  O  from  2.45  grms. 
of  KC1O3?  3.  How  much  HgO  is  necessary  to  yield  i  cu.  d.  m.  of  O  ? 
4.  How  many  times  is  O  heavier  than  H  ? 

OZONE.— O3. 

(i)  Pour  water  on  clean  pieces  of  phosphorus  to  half  cover  them  ; 
invert  a  large,  clean  jar  over  this  and  allow  to  stand  for  several  hours'. 
Test  the  air  under  the  jar  for  ozone.  For  this  purpose  use  paper  impreg- 
nated with  a  mixture  of  starch  paste  and  potassium  iodide.  What  occurs  ? 

(Read  Richter,  p.  85-89:) 

COMPOUNDS  OF  OXYGEN  AND  HYDROGEN. 

WATER.— H2O. 

(1)  Arrange  the  distillation  apparatus  (Fig.  17)  and  prepare  about  100 
cc.  of  distilled  water.     Note  its  taste  and  odor.     Test  it  for  chlorides 
with  AgNO3.     Does  it  leave  a  residue  FlG.  I7. 

upon  evaporation  ?     What  action  has 
it  on  litmus? 

Apply  all  these  tests  to  a  natural 
water  (except  rain). 

(2)  i.  Heat  a  little  vegetable  mat- 
ter in  a  dry  test-tube.      2.  Heat  fresh 
meat  in  the  same  manner.     3.   Care- 
fully heat  crystals  of  zinc  or  copper 

sulphate  in  a  test-tube.  What  happens  in  these  experiments?  4.  Expose 
clear  crystals  of  sodium  phosphate,  on  a  watch-crystal  to  the  air.  5.  Do 
the  same  with  pieces  of  calcium  chloride.  Results  ? 

(3)  Determine  the  quantitative  composition  of  water. 

1.  The  composition  of  water  by  weight  follows  from  the  experiment  of 
reducing  oxide  of  copper  described  under  H. 

2.  The  relative  volumes  with  which  O  and  H  unite  to  form  water,  are 
determined  either  by  analysis  or  synthesis.     The  former  has  been  per- 
formed in  electrolyzing  water. 

3.  Fill  a  eudiometer  (Fig.  18)  with  water.     Through  a  rubber  tube 
admit  about  50  cc.  of  O  and  then  a  like  volume  of  H.     (If  the  eudiome- 
ter is  not  graduated,  mark  these  with  rubber  bands.)     Close  the  open 
end  with  your  thumb,  leaving  some  air  to  serve  as  a  cushion  beneath  it, 
and  pass  the  spark.     Remove  the  thumb,  and  pour  in  enough  water  to 


22 


EXPERIMENTS    IN    GENERAL   CHEMISTRY. 


FIG.  18. 


make  the  levels  equal  in  both  limbs.     What  is  the  amount  of  the  contrac- 
tion ?    What  is  the  residual  gas  ?     Test  it. 

(4)  Determine  the  weight  of  a  litre  of  steam, — Con- 
struct apparatus  shown  in  Fig.  19.     The  flask  a  is  closed 
with  a  cork.      C  is  a  vessel  containing  melted  paraffin. 
A  small  glass  tube  is  weighed  and  filled  with  water  (not 
more  than  .02  grm.).     Heat  the  vessel  Cwith  a  Bunsen 
burner  until  the  temperature  of  a  is  constant  (?).     Now 
drop  the  tube  containing  the  water  through  the  mouth 
of  the  flask  (the  bottom  of  which  should  be  protected 
with  a  layer  of  asbestos)  and  quickly  re-cork.     When 
the  fall  of  water  in  the  graduated  tube  ceases,  read  the  volume  of  gas, 
and  note  the  temperature  and  pressure  of  the  air. 
The  calculation  is  analogous  to  that  used  under  O. 
(5)  Perform  experiment  2,  p.  100  in  Richter. 
How  many  volumes  of  steam  result  from  the  com- 
bination of  2  vols.  of  H  and  i  vol.  of  O? 

How  would  you  deduce  the  molecular  formula  of 
water  from  the  preceding  experiments  ? 


FIG.  19. 


(1)  Add  moist  hydrated  barium  peroxide  to  cold  dilute  H2SO4.     Filter. 
What  does  the  filtrate  contain  ? 

(2)  i.  Add  a  solution  of  KI,  containing  starch,  to  a  portion  of  this 
liquid  (?).     Ferrous  sulphate  hastens  the  reaction.     2.   Cautiously  add  a 
dilute  solution  of  potassium  permanganate  to  another  portion  (?). 


COMPOUNDS  OF  OXYGEN  AND  CHLORINE. 

(1)  Make  a  dilute  solution  of  caustic  potash,  and  conduct  chlorine  into 
it  until  the  latter  is  no  longer  absorbed.     Treat  one  portion  of  the  product 
with  HC1,  and  another  with  H2SO4.     What  results  ? 

(2)  Mix  10  grms.  of  quicklime  with  25  cc  of  H2O.     After  the  slaking 
FIG.  20.  is  finished,  conduct  Cl  into  the  mixture  until  it  is  no 

longer  absorbed. 

Add  HC1  to  one  portion  and  H-jSO*  to  a  second 
portion. 

What  is  set  free  ?     Does  it  bleach  ? 

(3)  Pass  Cl  into  a  hot  concentrated  solution  of  KOH 
till  it  ceases  to  be  absorbed  (Fig.  20).  What  separates 
upon  cooling?  Recrystallize  the  product  from  water.  Will  it  give  off  O 


SECOND    NATURAL    GROUP   OF    ELEMENTS — SULPHUR.  23 

upon  heating?     Try  the  action  of  HC1  upon  a  crystal.     Allow  a  drop  of 
cone.  H-jSOi  to  fall  upon  a  small  crystal  and  warm  gently  (?).     Care  ! 
Observe  carefully  the  behavior  of  KC1O3  upon  heating  (?). 

SULPHUR.— S. 

(i)  Place  a  few  grams  of  powdered   S  in  a  dry  test-tube,  and  heat 
gradually.     Observe  and  describe  the  changes  which  occur. 

(3)  Dissolve  a  little  S  in  CS2  and  allow  to  stand   till  the  liquid  has 
evaporated.     What  remains? 

(4)  Determine  the  sp.  gr.  of  S  (Fig.  21).    Water,  previously  boiled  is 
introduced  into  a  flask.     It  is  essential  that  the  neck  of  the  flask  should 
be  narrow.     Weigh  the  flask,  then  place  an  additional  10  grm. 
weight  upon  the  right-hand  pan  of  the  balance  and  small  pieces     FIG.  21. 
of  S  upon  the  left-hand  pan,  until  the  pointer  is  again  in  the 
middle.     Now  introduce  the  S  into  the  flask.     Carefully  remove 

water  above  the  mark  and  re-weigh  the  flask  with  its  contents. 
The  loss  in  weight  will  represent  the  weight  of  a  volume  of  water 
equal  to  that  of  10  grms  of  S.  The  latter  divided  by  the  former 
is  the  specific  gravity  of  the  sulphur. 

(5)  Prepare  the  monoclinic  modification  of  S  by  melting  about  10  grms. 
of  the  ordinary  variety  in  a  covered  Hessian  crucible.     Cool ;  and  as 
soon  as  a  solid  crust  has  formed  upon  the  surface,  pierce  it  and  allow  the 
still  liquid  portion  of  the  contents  to  run  out.     Note  the  shape  of  the 
crystals  upon  the  sides  of  the  crucible. 

(6)  To  obtain   the  plastic  variety,  heat  10  grams  of  S  in  a  test-tube 
above  230°  C.,  and  pour  the  mass  into  cold  water. 

Test  the  solubility  of  the  product  in  CS2.     Preserve  a  portion  of  it  for 
several  days.     Does  it  change  ? 

(7)  To  a  strong   solution  of  yellow    potassium   sulphide,  add    HC1. 
What  are  the  properties  of  the  separated  sulphur  ? 

Give  a  brief  outline  of  the  element  sulphur ;  compare  it  with  the  pre- 
viously studied  elements. 

SULPHUR   AND    HYDROGEN. 

(8)  Hydrogen  sulphide  is  formed  with  difficulty  from  its  elements,  but 
is  readily  obtained  by  the  action  of  acids  upon  sulphides,  thus  : — 

FeS  +  H2SO4  =  FeSO4  +  H2S  or  Sb2S3  +  6HC1  =•  2SbCl3  +  3H2S. 

The  apparatus  to  be  used  is  the  same  as  that  employed  in  preparing 
hydrogen. 

(9)  What  are  the  properties  of  H2S  ?     Is  it  soluble  in  water  ?     Does  it 


24 


EXPERIMENTS    IN    GENERAL   CHEMISTRY. 


burn  ?  What  are  its  products  of  combustion  ?  Hold  a  porcelain  plate  in 
the  flame;  what  results  ?  Pass  the  gas  into  solutions  of  chromic  acid,  per- 
manganic acid,  and  ferric  chloride.  What  changes  are  observed  ?  How 
do  these  last-named  reactions  show  its  reducing  power?  What  happens. 
to  the  aqueous  solution  of  the  gas  when  exposed  to  the  air  ?  What  action 
has  the  aqueous  solution  upon  litmus?  To  what  class  of  compounds  does 
it,  therefore,  belong? 

(10)  Pass  H2S  through  solutions  of  the  following  salts,  viz.  :  —  CuSO4, 
SbCl3,  Pb  (NO3)2,  AsCl3,  and  Zn  (C2H3O2)2.     Note  results  carefully. 
Can  sulphides  be  prepared  in  another  manner?     (See  Chap.  II,  §  i.) 
(n)  Determine  the  composition  of  hydrogen  sulphide. 

Into  a  bent  tube  of  hard  glass,  filled  with  mercury, 
(Fig.  22),  introduce  dry  hydrogen  sulphide.*     Place  a 
piece  of  tin  in  the  bent  portion,  and  heat   it.     Is  the 
volume  of  the  gas  changed  after  the  experiment,  and 
what  becomes  of  the  piece  of  tin  ?     Test  the  gas  re- 
maining in   the  tube.     Do  your  results  enable  you  to 
deduce  the  molecular  formula  of  H2S  ?     (See  Richter, 
3d  ed.,  p.  in.)     Trace  the  similarity  between  H2S  and   H2O.     Write  a 
summary  of  your  experiments  on  H2S. 


(12) 


FIG.  23. 


SULPHUR    AND    CHLORINE. 

Sulphur  Mono  chloride. — Prepare  this  compound  by  conducting 
dry  chlorine  over  molten  sulphur.  The  product  which  distils 
over  is  collected  in  a  dry  test-tube,  kept  cold  by  immersion 
in  ice  water.  2.  Redistil  the  product.  Determine  its  boiling 
point  in  an  apparatus  similar  to  that  pictured  in  Fig.  23. 
Note  the  color  and  odor  of  the  product.  Expose  some  of  it 
to  the  air  on  a  watch-glass.  Add  water  to  another  portion 
contained  in  a  test-tube.  Note  carefully  what  happens. 
Write  the  reaction,  and  examine  for  all  the  products. 


SULPHUR    AND    OXYGEN. 

(13)  Burn  sulphur  in  the  air.     Result?     Burn  FeS2  in  the  air.     What 
are  the  properties  of  the  resulting  compound  ?     It  is  sulphur  dioxide  —  SO2. 

(14)  Fit  a  small  flask,  as  indicated  in  Fig.  24.     Place  copper  turnings 
in  it,  then  add  H2SO4  (strong)  through  the  funnel  tube.     Warm.     Is  the 
product  the  same  as  that  obtained  in  13?     Is  it  soluble  in  water?     Has 
the  aqueous  solution  the  same  properties  as  the  gas?     2.   Pass  some  of 


The  instructor  should  perform  this  experiment. 


NITROGEN  GROUP NITROGEN.  25 

the  gas  into  solutions  of  potassium  dichroraate  and  potassium  permanga- 
nate acidulated  with   H2SO4.     Repeat   these   experiments          Fir,.  24 
with    the   aqueous    solution    instead    of  the    gas.     What 
happens  in  each  case?    3.  Test  the  aqueous  solution  of 
SO2  with  litmus.     What  is  this  solution  commonly  called  ? 
4.   Fill  a  dry  jar  with  SO2  gas  ;  introduce  colored  flowers. 
Note  the  result. 

(15)  What   is   the  formula  of  sulphurous   acid?     How 
many  series  of  salts  can  it  form  ?     How  would  you  desig- 
nate the  different  sodium  salts  ?     Add  HC1  to  a  solution 
of  Na2SO3.     What    follows?     Evaporate    the   solution   to 
dryness  and  examine  the  residue.     What  is  it?    Write  the  reaction. 

SULPHUR  TRIOXIDE — SO3 — and  SULPHURIC  ACID — H2SO4.  (Read    Rich- 
ter,  p.  189). 

(16)  i.  Prepare  sulphuric  acid  as  described  in  Richter,  p.  191.     Study 
the  product  carefully.   2.   Dilute  a  portion  of  it  with  water  ;  what  happens  ? 
2.  Test  a  portion  of  this  diluted  solution  with  litmus  (?).     4.  Another 
portion   neutralize  with  NaOH  and  evaporate.     What   is  the  residue  ? 
Does  it  contain  any  S?     Prove  this.     5.  Add  BaCl2  to  a  third  portion  of 
the  solution.     What  is  the  precipitate  ?     Is  it  soluble  in  water  or  in  hydro- 
chloric acid  ?     6.  What  is  the  action  of  strong  H2SO4  upon  wood  or 
paper?     Explain  the  cause  of  this  action. 

(17)  How  many  series  of  salts  can   sulphuric   acid    form.       Prepare 
(NH4)2SO4,  Na2SO4,  NaHSO4and  CuSO4.     (Read  Richter,  pp.  189-200). 


CHAPTER  VI. 

NITROGEN  GROUP— NITROGEN,  PHOSPHORUS,  ARSENIC,  ANTIMONY 
AND  BISMUTH. 

NITROGEN.— N. 

(i)  Preparation. — i.   In  a  dish  swimming  on  water  place  a  piece  of 
phosphorus  and  ignite  it ;  invert  a  beaker  glass  over  it  (Fig.  25).     What 
FIG  25.          FIG  26         becomes  of  the  P  ?     When  the  latter  has  ceased 
burning,  restore  the  level  of  the  water,  and  note 
the  decrease  in  the  volume  of  the  air.     Test  the 
residual  gas  with  a  burn  ing  taper.     2.   Heat  gently 
in  a  small  flask  or  retort  a  mixture  of  i  part  KNO2, 
i   pt.  NH4C1,  i   pt.  K2Cr2O7,  and  3  pts.  of  H2O; 
collect  the  gas  over  water.     Fill  five  bottles  with 
this  gas. 

(2)  Has  it  color,  taste,  odor?     Does  it  burn  or 
4 


26 


EXPERIMENTS    IN    GENERAL   CHEMISTRY. 


FIG. 


FIG.  28. 


support  combustion  ?     Is  the  gas  heavier  than  air  ?     Does  it  unite  readily 
with  other  elements? 

(3)  Determine  the  weight  of  a  litre  of  nitrogen. — A  round-bottomed 
flask  is  fitted,  as  shown  in  Fig.  26.  Pour  about  30  cc.  of  water  into  it, 
and  insert  the  rubber  cork  to  the  mark.  Boil  the 
water,  while  the  clip  is  open,  until  all  the  air  has  been 
expelled  from  the  flask.  Steam  should  be  allowed  to 
escape  for  about  five  minutes.  Now  close  the  tube 
with  the  clip,  and  remove  the  flame.  Cool  and 
weigh  the  flask.  Read  the  temperature  and  baro- 
metric pressure  in  the  balance-room. 

Connect  the  flask  with  the  tube,  b,  of  the  aspirator, 
containing  N,   and    arranged  as  in   Fig.    27.      The 
rubber  tube,  a,  is  made  to  dip  under  water,  and  the 
clip  is  gradually  opened,  allowing  N  to  enter  the  flask.     Now  raise  the 
vessel  containing   the    water  into  which  the   rubber  tube 
dips,  so  that  the  water  in  it  is  at  a  higher  level  than  that 
in  the  aspirator.     Close  the  clip.     Disconnect  the  flask 
and  open  the  clip  for  a  moment,  to  establish  atmospheric 
pressure  in  the  flask.     Weigh.    The  calculation  is  identical 
with  that  given  for  O. 

Wnat  is  the  ratio  between  the  weights  of  equal  volumes 
of  N  and  H? 

(4)  Is  air  a  chemical  compound  ? 
How  would  you  determine  the  weight  of  a  litre  of  air  ? 
(5)   i.  Determination  of  the  Oxygen  in  Air  by  the  Pyrogallate  Method, — 
At    the    atmospheric    temperature    and    pressure 
measure  off  50  cc.  of  air  in  the  Hempel  burette, 
shown   in  Fig.    28.     Connect  this  at   c  with  the 
capillary  of  a  Hempel's  compound  pipette  (Fig. 
29)  containing  an  alkaline  solution  of  pyrogallate 
of  potash.     Open  the  clips  and  transfer  the  air  to 
the  pipette  by  raising  the  tube,  a.     When  this  is 
accomplished,  and  the  capillary  of  the  pipette  is  filled  with  water  from  b, 
close  the  clips  again.    Disconnect  the  apparatus.    Shake 
the  pipette  for  several  minutes  so  as  to  bring  gas  and 
absorbent  in  intimate  contact.      Reconnect  pipette  and 
burette,    and    force    the    residual    gas   into   the   latter. 
Restore    atmospheric    pressure   and    read    the   volume. 
What  does  the  loss  represent  ? 

2.  Explosion  Method. — To  40  cc.  of  air  contained  in 


FIG.  29. 


NITROGEN    GROUP — AMMONIA.  27 

the  burette  add  40  cc.  of  pure  H.  Pass  this  mixture  into  the  Hempel 
explosion  pipette  shown  in  Fig.  30.  Close  the  stop-cock,  d,  and  the  clip, 
c,  then  connect  the  platinum  electrodes  with  an  inductor  and  pass  a  spark 
What  takes  place  ?  Measure  the  volume  of  the  gas  remaining.  How 
much  of  the  contraction  was  due  to  O  ?  What  is  the  composition  of  the 
gas  after  the  explosion  ? 

(Study  Richter,  pp.  116-125.) 

NITROGEN  AND  HYDROGEN. 
AMMONIA. 

(6)  Preparation. — Heat  an  intimate  mixture  of  finely  powdered  ammo- 
nium  chloride  and   caustic  lime   in  a  flask ;    conduct   the    evolved    gas 
through  a  tube  filled  with  small  pieces  of  lime,  and  collect  it  in  jars  or 
test-tubes  over  mercury. 

What  is  the  object  of  the  lime  in  the  tube?  Why  can  you  not  dry 
the  gas  by  passing  it  through  H2SO4  or  CaCl2  ?  Why  should  it  be  col- 
lected over  mercury  ? 

(7)  Is    ammonia  gas    combustible?     Does   it  FIG.  3i.          FIG.  32. 
support  combustion  ?     i.   Conduct  NH3  through 

a  glass  tube,  and  insert  this  into  a  wider  tube 
filled  with  oxygen  (Fig.  31).  Apply  a  flame. 
The  ammonia  gas  will  ignite  and  continue  to 
burn.  2.  Heat  concentrated  ammonia  water  in 
a  beaker  until  there  is  an  abundant  disengage- 
ment of  gas,  then  conduct  a  rapid  current  of 
oxygen  through  the  liquid,  and  lower  a  glowing 
spiral  of  platinum  into  the  beaker  (as  in  Fig.  32). 
What  happens  ? 

Note  the  odor  of  NH3  (caution  ?).  Is  it  lighter  than  air?  Soluble  in 
water  ? 

(8)  Prepare  an  aqueous  solution  of  ammonia. 
What  are  its  properties  ? 

Add  red  litmus  to  some  of  the  solution  (?),  and  then  neutralize  care- 
fully with  dilute  HC1.  Evaporate  to  dryness.  Compare  the  product 
with  the  ordinary  NH4C1.  Test  it  for  Cl  (?).  Heat  a  little  of  it  with 
sodium  hydroxide  (?).  Heat  another  portion  on  a  platinum  foil  (?). 

(9)  Determine  the  weight  of  a  litre  of  NH^ 

Fill  a  dry  flask  with  the  gas  by  upward  displacement,  and  proceed 
exactly  as  under  chlorine.  What  is  the  density  of  NHS  ? 

To  determine  the  quantitative  composition  of  NH3,  perform  experi- 
ments T  and  2  on  pp.  130  and  131,  in  Richter. 

Write  out  summary.     (Read  Richter,  pp.  125-131.) 


28  EXPERIMENTS    IN    GENERAL    CHEMISTRY. 

NITROGEN  AND  THE  HALOGENS. 

(n)  Pour  a  saturated  alcoholic  solution  of  iodine  into  strong  ammonia 
Weuer.  Collect  the  precipitate  on  a  filter  and  wash  it  with  water.  Open 
the  moist  filter;  tear  it  into  small  pieces  and  spread  these  on  a  board. 
After  they  have  become  dry,  touch  them  with  the  end  of  a  rod  (?).  Ask 
for  instructions  /  (Read  Richter,  pp.  132-133.) 

NITROGEN  AND  OXYGEN. 

(12)  Hyponitrous  oxide — N2O.      i.   Place  about  5  grms.  of  ammonium 
nitrate  in  a  small  retort ;  add  a  little  water,  and  apply  heat.     Collect  the 
product  over  warm  water.      2.  Test  it  with  a  glimmering  chip ;  3.   with 
burning  phosphorus ;  4.   with  burning  sulphur.     5.   Mix  equal  volumes  of 
this  gas  and  of  H,  and  apply  a  flame.     What  other  gas  does  it  resemble 
in  its  properties?     (Read  Richter,  pp.  212-213.) 

(13)  Nitric  oxide — NO.      i.    Pour  dilute  HN03   (sp.    gr.    1.2)  upon 
copper  turnings  contained   in  an   evolution   flask.     Cool,  and  allow  the 
red  fumes,  which  form  at  first,  to  escape  ;  then  collect  the  colorless  pro- 
duct over  water.      2.   What  occurs  when  this  gas  comes  in  contact  with 
the  air?     Is  it  the  O  or  the  N  of  the  air  that  acts  upon  the  gas?     3. 
Apply  the  tests  given  under  (12)  to  this  gas  (?).     How  can  NO  be  dis- 
tinguished from  oxygen  ?     4.   Fill  a  cylinder  with  NO,  and  add  a  few 
drops  of  CS2,  shake  well  and  bring  a  flame  to  the  mouth  of  the  vessel  (?). 
5.   Pass  a  current  of  NO  into  a  strong  solution  of  ferrous  sulphate.     What 
occurs?     After  the  solution  has  become  saturated  with  the  gas  heat  it  to 
boiling  (?).     6.   Pass  the  gas  into  a  solution  of  potassium  permanganate  (?). 

(14)  Nitrogen  trioxide — N2O8.     (Read  Richter,  pp.  205-206.) 
Nitrous  Add—  HNO2.     (Richter,  p.  206.) 

(15)  Nitrogen  tetroxide,  N2O4,  and  dioxide,  NO2.      i.  Heat  10  grams 
of  dry  lead  nitrate  in  a  test  tube  ;  condense  the  escaping  vapors  in  a 
well-cooled  receiver.     What  are  the  vapors,  and  what  is  the  condensed 
liquid  ?     Note  the  color.     2.   What  is  the  action  of  cold  water,  and  of 
aqueous  solutions  of  the  alkalies  upon  N2O4?     What  do  these  reactions 
indicate  in  respect  to  the  composition  of  this  compound  ?     (Richter,  pp. 
207-208.)     3.  What  is  its  action  upon  potassium  iodide? 

(16)  Nitrogen pentoxide,  N2O5.     (Richter,  p.  205.) 

NITRIC  ACID.— HNO3. 

i.  Preparation. — In  a  retort  heat  a  mixture  of  sodium  nitrate  and  sul- 
phuric acid  in  proportions  corresponding  to  the  equation  (?)  : 
NaNO,  +  H2SO4  =  N*aHSO4  -f  HNO3. 


NITROGEN    GROUP — PHOSPHORUS    AND    HYDROGEN.  29 

Collect  the  product  in  a  cold  receiver. 

2.  What  are  the  physical  properties  of  HNO3  ?  Color  ?  Odor  ? 
Action  on  litmus  (dilute  with  H2O)  ?  3.  What  action  has  it  on  indigo  ? 
Upon  the  skin  ?  4.  Notice  the  effect  of  the  acid  upon  the  following 
metals:  Cu,  Fe,  Pb,  Zn,  Sn.  Write  the  reaction  for  each  one.  5. 
Cover  powdered  sulphur  with  the  acid,  and  warm  (?).  Dilute  with  water, 
filter,  and  test  the  liquid  with  BaCl2  (?).  6.  Add  a  few  drops  of  HNO3  to 
a  solution  of  ferrous  sulphate  (?) ;  warm  the  solution  (?). 

Problems. — i.  Required  i  cu.  m.  of  N.  How  much  air  is  to  be  de- 
prived of  O ;  and  how  much  P  must  be  burned,  if  62  pts.  of  P  unite 
with  80  pts.  of  O  ? 

2.  How  much  HNO3,  containing  46  per  cent,  of  water,  may  be  ob- 
tained from   1,700   grms.  of  NaNO3,  and  how  much  water  must  be  taken  ? 

3.  How  many  grams  of  NH3  will  be  absorbed  by  5  litres  of  H2O,  if 
the  latter  absorbs  500  times  its  volume  of  the  gas  ?     4.  Ten  litres  of 
water  having  absorbed  700  times  their  volume  of  ammonia,  what  are  the 
least  amounts  of  NH4C1  and  CaO  necessary  for  producing  this  solution  ? 

*  PHOSPHORUS.— P. 

(1)  i.  Determine  the   physical  properties  of  the  active  and  the  red 
varieties.     2.  Allow  a  small  piece  of  the  active  variety  to  ignite  in  the  air. 
Will   the  red  variety  do   this?     3.   Throw  a  small  piece  of  the  yellow 
variety  into   a  jar  of  dry   Cl    (?).      Repeat  with   the   red    variety    (?). 
4.  Bring  a  small   dry  piece  of  active   P  in  contact  with  iodine  (?).     5. 
Heat  a  flask  containing  a  small  piece  of  P  and  water  until  the  former  is 
melted,  then  pass  a  current  of  oxygen  through  a  delivery  tube  into  the 
melted  phosphorus  (?).      Care  !     (Study  Richter,  pp.  133-136.) 

PHOSPHORUS  AND  HYDROGEN. 

(2)  Phosphine — PH3.      i.    Fill  a  flask  almost  full   with   a  moderately 
concentrated  NaOH  solution.     Add  a  few  pieces  of  P,  and  heat  carefully. 
When  the  air  in  the  neck  of  the  flask  has  been  expelled  by  the  escaping 
gas,  insert  a  cork  with  a  delivery  tube  the  other  end  of  which  dips  under 
warm  water.     What  becomes  of  the  gas  as  it  escapes  into  the  air  ?    Write 
the  reaction  involved. 

(Richter,  pp.  136-139.) 

Is  there  any  similarity  between  PH3  and  NH3  ? 


30  EXPERIMENTS    IN    GENERAL    CHEMISTRY. 

PHOSPHORUS  AND  THE  HALOGENS. 

(3)  i.   Pass  a  current  of  dry  CO2  gas  into  a  retort,  the  bottom  of  which 
is  covered  with  dry  sand.     When  all  the  air  has  been  expelled,  introduce 
some  well-dried  pieces  of  P,  and  replace  the  CO2  by  a  stream  of  dry  Cl. 
Connect  the  neck  of  the  retort  with  a  Liebig's  condenser,  and  collect 
the  product  in  a  receiver.     It  is  phosphorus  trichloride.     What  are  its 
properties?     Pour  some  of  it  into  water  (?). 

2.  Place  a  little  PC13  in  a  dry  test-tube,  and  pass  a  stream  of  dry  Cl 
upon  its  surface.  What  is  the  result  ? 

PHOSPHORUS  AND  OXYGEN. 

(Richter,  pp.  214-219.) 

(4)  i.   Prepare  phosphorus pentoxifte,  P2O5,  by  burning  a  carefully  dried 
piece  of  P  under  a  dry  bell -jar.      2.   Drop  a  portion  of  the  product  into 
water  (?). 

(5)  Orthophosphoric   acid,    H3PO4 ;    metaphosphoric  acid,  HPO3;   and 
pyrophosphoric  acid,  H4P2O7.     How  are  these  acids  obtained  ?     How  many 
series  of  salts  are  derived  from  them  ?     By  what  names  would  y<Ju  distin- 
guish the  different  salts  ? 

i.  Dissolve  some  Na2HPO4  in  water  and  test  che  solution  with  AgNO3, 
and  FeCl3.  What  do  you  observe  in  each  case  ?  2.  Dissolve  fused 
Na2HPO4  in  water,  and  perform  the  same  tests  with  its  solution.  3.  Heat 
salt  of  phosphorus  (NaNH4HPO4)  until  it  no  longer  effervesces  ;  cool, 
crush  the  residue  in  a  mortar,  and  dissolve  it  in  water.  How  does  this 
solution  behave  upon  treating  with  the  above  reagents?  4.  Acidify  a 
portion  of  the  last-named  solution  with  acetic  acid,  and  add  a  solution  of 
albumen  to  it.  Result? 

(6)  Phosphorus  trioxide — P2O3,  and  phosphorous  acid — H3PO3. 

Pour  PC13  into  water.     Evaporate  the  solution  to  syrupy  consistency  (?). 
(Study  Richter,  p.  216.) 

(7)  Hypophosphorous  acid — H3PO2. 

Heat  pieces  of  phosphorus  in  a  porcelain  dish  with  a  moderately 
strong  baryta  solution  (see  p.  29).  When- no  more  PH3  is  formed,  cool, 
filter,  and  pass  CO2  into  the  solution  until  it  shows  a  neutral  reaction  to 
litmus.  Toward  the  end,  the  solution  should  be  warmed.  Filter  and 
evaporate  to  suitable  concentration.  Hypophosphite  of  barium  will 
crystallize. 

How  may  the  free  acid  be  obtained  from  this  salt  ? 


NITROGEN   GROUP — ANTIMONY.  31 

ARSENIC.— As. 

(1)  Study   the   physical  and   chemical    properties    of   this    element. 
(Richter,  pp.  142  and  143.)     Are  they  analogous  to  those  of  phosphorus  ? 

i.  In  a  tube  of  hard  glass  heat  a  small  piece  of  As  to  redness.  Result  ? 
2.  Heat  As  with  the  oxidizing  flame  upon  charcoal  (?).  3.  Dissolve 
powdered  As  in  strong  HNO3  (?). 

ARSENIC  AND  HYDROGEN.  * 

(2)  Perform  Marsh' s  test  for  As* 
Arrange  the  apparatus  shown  in  Fig.    ^ 

33.  To  the  mixture  of  Zn  and  dilute 
H2SO4  contained  in  a,  add  a  small  por- 
tion of  the  solution  to  be  tested  for  As. 
The  liberated  gas  contains  H  and  AsH3 
(arsine).  It  is  passed  through  c,  filled 
with  CaCl2  (?),  and  then  through  d,  a 
tube  of  hard  glass,  contracted  at  several  places.  After  all  the  air  has 
been  expelled  from  the  apparatus,  ignite  the  hydrogen.  If  As  is  present 
it  will  burn  with  a  bluish  white  flame,  and  white  vapors  will  be  given  off. 
Hold  a  cold  porcelain  plate  in  the  flame  (?).  Heat  the  tube  d,  as  shown 
in  the  figure  (?). 

Great  care  must  be  exercised  in  performing  this  test,  as  the  arsine  gas  is 
extremely  poisonous  ! 

ANTIMONY.— Sb. 

(i)  Study  this  element  in  the  same  manner  as  As.  Distinguish  between 
SbH3  and  AsH3. 

i.  Treat  the  metallic  mirrors  obtained  in  Marsh's  apparatus,  with  a 
freshly  prepared  solution  of  hypochlorite  of  sodium  :  As  dissolves  readily, 
while  Sb  is  scarcely  acted  upon.  2.  Heat  a  piece  of  the  tube  in  which  a 
mirror  has  formed,  in  the  flame  of  the  Bunsen  burner.  Dissolve  the  pro- 
duct in  dilute,  warm  HCI,  and  add  H2S  water  (?).  3.  Treat  the  spot 
formed  upon  a  cold  porcelain  plate  with  yellow  ammonium  sulphide,  and 
evaporate  the  solution  at  a  gentle  heat  (?). 

Problems. — (i)  How  much  P  can  be  obtained  from  250  grms.  of 
bones?  (See  Richter,  p.  134.)  (2)  10  grms.  of  P  give  what  vol.  phos- 
phine  ?  (3)  What  is  the  weight  of  the  product  remaining,  after  evapor- 
ating a  solution  of  10  grms.  of  As  in  HNO3? 

*  Ask  for  instructions. 


32  EXPERIMENTS    IN    GENERAL    CHEMISTRY. 

CHAPTER  VI. 

CARBON  GROUP-CARBON  AND  SILICON. 
CARBON.— C. 

(i)  How  many  allotropic  modifications  of  this  element  are  known? 
What  are  their  principal  properties?     i.  Boil  a  dilute  litmus   solution 
with  powdered  animal  charcoal;  filter.     Result?     2.  Substitute  indigo 
for  the  litmus  in  the  preceding  experiment  (?). 
3.  Determination  of  the  composition  of  coal. 

i.  Volatile  matter  and  coke.  Weigh  out  2  grms.  of  powdered  coal  in 
a  platinum  crucible  provided  with  a  well-fitting 
cover.  Heat  with  a  large  flame,  until  the  escaping 
gases  cease  to  burn  between  the  lid  and  the 
crucible.  A  blast  lamp  flame  is  applied  for  a 
minute  longer.  Cool  and  weigh.  Loss  in  weight 
represents  the  volatile  matter.  The  residue  is 
called  coke.  , 

2.  Ash. — A  second  portion  of  coat  (i  grm.)  is 
gently  heated   over  the  Bunsen   flame,  until  the 
volatile  constituents   are  expelled.     The   heat  is 
then  raised  and  the  lid  of  the  crucible  placed  in  the  position  indicated  in 
Fig.  34.     The  residue  is  the  ash. 
(Read  Richter,  pp.  150-152.) 


CARBON  AND  HYDROGEN. 

(2)  Methane  (Marsh  gas)— CH4. 

i.  Preparation. — Heat  a  dried  mixture  of  sodium  acetate  and  sodium 
hydroxide  in  an  iron  tube.*  Collect  the  gas  over  water.  Note  its  color, 
odor  and  taste.  Does  it  burn?  2.  Mix  i  vol.  of  it  with  7  to  8 
times  its  vol.  of  air  and  explode  by  applying  a  flame.  (Ask  for  in- 
structions !) 

How  would  you  determine  the  molecular  weight  of  this  compound  ? 

(3)  Make  a  eudiometric  combustion  of  i  vol.  of  CH4  with  2  vols.  of 
O  as  described  in  Richter,  p.  121. 

(4)  Ethane— C2H6.     (Richter,  p.  153.) 

(5)  Acetylene — C2H2.     Light  a  Bunsen  burner  at  the  base  and  turn  it 

*A  hard  glass  tube  will  answer. 


CARBON    GROUP — SILICON.  33 

down,  so  that  the  flame  is  small.     Acetylene  can  be  recognized,  among 
the  products  of  combustion,  by  its  characteristic  odor. 

(6)  CARBON  AND  THE  HALOGENS.     (Richter,  p.  160.) 

CARBON  AND   OXYGEN. 

( 7 )  Carbon ,  dioxide — CQ2. 

i.  Preparation. — Upon  pieces  of  marble,  contained  in  an  evolution 
flask,  pour  dilute  HC1  (i  HC1  :  1-2  H2O).  Conduct  the  resulting  gas 
through  water  and  through  cone.  H.jSO4.  It  may  be  collected  either  by 
downward  displacement  of  the  air,  or  over  mercury.  2.  Note  color, 
taste  and  odor  of  this  gas.  Is  it  soluble  in  water  ?  How  does  its  weight 
compare  with  that  of  air  ?  Does  it  burn  or  support  combustion  ? 
3.  Conduct  a  current  of  CO2  into  a  solution  of  NaOH,  evaporate  the 
liquid,  and  test  the  residue  for  Na2CO3(?).  4.  To  different  portions  ,of 
Na2CO3  solution,  add  solutions  of  MgSO4i  BaCl2>  Pb  (NO3)2,  ZnSO4  (?). 

(Study  Richter,  pp.  228-232.) 

(8)  Carbon  monoxide — CO. 

Preparation. — i.  In  a  tube  of  hard  glass  heat  zinc  dust  to  faint  redness, 
while  conducting  a  slow  current  of  CO2  over  it.  In  what  respect  does 
the  product  differ  from  CO2.  2.  Heat  crystals  of  oxalic  acid  with 
cone.  H2SO4  in  a  flask,  and  wash  the  product  with  a  NaOH  solution. 
Write  the  reaction.  Study  the  properties  of  this  gas.  (Richter,  p.  233.) 

(9)  Carbon  disulphide — CS2. 

Perform  some  of  the  experiments  indicated  in  Richter,  p.  234. 

(10)  CARBON  AND  NITROGEN. 

i.  In  a  dry  test-tube  heat  a  nitrogenous  carbon  compound  with  a  small 
piece  of  K.  Cool  and  add  water.  KCN  is  formed  and  can  be  tested 
with  AgNO3.  2.  Convert  a  portion  of  the  KCN  into  KCNS  by  evapo- 
rating with  (NH4)2S.  Test  with  FeCl3.  3.  To  a  solution  of  FeSO4  add 
potassium  ferrocyanide.  What  results  ?  4.  What  is  the  action  of  the 
ferrocyanide  upon  solutions  of  ferric  salts  ? 

(n)  Study  the  nature  of  flame.  Make  the  experiments  described'in 
Richter,  pp.  155-160. 

SILICON.— Si. 

(i)  Preparation. — Make  an  intimate  mixture1  of  i  grm.  magnesium 
powder  and  4  grms.  of  finely  powdered  quartz-sand.  Heat  this  to 
bright  redness  in  a  wide  tube  of  hard  glass.  It  is  best  to  use  the  blast 
lamp  for  this  purpose.  The  part  of  the  tube  containing  the  mixture 
should  be  rotated  in  the  flame.  The  residue,  after  a  few  minutes'  heating, 
is  allowed  to  cool,  and  treated  with  water  containing  HC1.  The  product 
5 


34  EXPERIMENTS    IN    GENERAL    CHEMISTRY. 

consists  of  amorphous  silicon  and  undecomposed  quartz.  2.  Test  the 
action  of  the  following  reagents  upon  Si  :  sulphuric,  nitric  and  hydro- 
fluoric acids,  potash  solution  and  chlorine.  (Read  Richter,  p.  161.) 


SILICON  AND  OXYGEN. 

(2)  Silicon  dioxide  (Silica,  Quartz) — SiO2. 

i.  Test  its  solubility  in  the  various  acids  and  alkalies.  2.  Fuse  a  mix- 
ture of  i  grm.  of  finely  powdered  quartz  with  4  grms.  of  Na.2CO3,  in  a 
platinum  crucible.  Dissolve  the  product  in  water.  3.  To  a  portion  of 
this  solution  add  HC1,  and  evaporate  to  complete  dryness.  Take  up  the 
residue  with  water  and  filter  off  the  insoluble  portion.  4.  To  another 
portion  of  the  aqueous  solution  of  the  fusion  add  NH4C1.  (?).  Make  a 
bead  of  salt  of  phosphorus  ;  bring  a  fragment  of  a  silicate  or  of  quartz  into 
it,  and  heat  in  the  blow-pipe  flame  for  a  few  minutes  (?). 

BORON.— B. 

(1)  Preparation  similar  to  that  of  Si.     What  are  its  properties?     Does 
it  unite  directly  with  other  elements  ?     Is  it  known  in  several  allotropic 
modifications  ?     What  is  the  valency  of  this  element  ? 

(Read  Richter,  pp.  240  and  241.) 

BORON  AND  OXYGEN. 

(2)  Boric  Acid— -BO3. 

i.  Dissolve  borax  in  5  parts  of  boiling  water,  add  HC1  to  acid  reaction, 
and  allow  to  cool.  What  crystallizes  out  of  the  solution  ?  Dry  some  of 
the  product  by  pressing  it  between  filter  paper.  Test  its  solubility  in 
water  and  in  alcohol.  What  do  you  observe  on  igniting  the  alcoholic 
solution  ?  Moisten  a  piece  of  turmeric  paper  with  an  aqueous  solution  of 
boric  acid,  and  dry  at  a  gentle  heat.  What  happens? 


Problems. — (i)  How  much  CO2  results  from  the  combustion  of  12 
grms.  of  carbon  ?  (2)  How  much  CO2  will  an  indefinite  quantity  of 
CaCO3  give,  when  acted  upon  by  4.666  grms.  of  muriatic  acid,  contain- 
ing 30  per  cent,  of  pure  HC1  ?  (3)  How  many  cubic  decimeters  of  CO 
can  be  obtained  from  90  grms.  of  oxalic  acid  ?  (4)  What  amount  of  SiO.2 
can  be  obtained  from  2  grms.  of  Wollastonite  (CaSiO3)  ?  (5)  What  is 
the  theoretical  quantity  of  boric  acid  obtainable  from  15  grms.  of  borax 
(Na2B4O7-f  ioH2O)? 


METALS    OF    THE    ALKALIES — POTASSIUM    AND    OXYGEN.  35 

METALS. 


CHAPTER  VII. 

METALS  OF  THE  ALKALIES— POTASSIUM,  SODIUM,  [AMMONIUM]. 

POTASSIUM.— K. 

(1)  Preparation. — Arrange  apparatus  as  shown  in  Fig.  35.     Into  a  tube 
of  hard  glass,  c,  introduce  a  porcelain  boat  containing  about  i  grm.  of  a 
mixture  of   138    pts.   (i  FJG  35 

mol.)  of  dry  (?)K2CO3 
and  72  pts.  (3  at.)  of 
Mg  powder.  Pass  a  cur- 
rent of  dry  H  over  it, 
and  after  all  the  air  has 
been  displaced  in  the 
apparatus  (?),  light  the 
escaping  gas ;  heat  the 
part  of  the  tube  surround- 
ing the  boat  to  incipient  redness.  Observe  the  brilliant  metallic  mirror 
which  is  formed,  and  drive  it  away  from  the  boat  by  increasing  the 
temperature  :  it  is  potassium.  Note  also  the  green  color  of  the  vapor 
and  the  violet  coloration  it  imparts  to  the  burning  hydrogen.  What  is  the 
residue  left  in  the  boat  ?  Test  its  reaction  with  litmus  (?). 

Formulate  the  reaction  involved  in  this  method  of  preparation. 

(2)  i.   Cut  a  piece  of  K  with  a  knife,  and  observe  the  color  and  lustre  of 
the  fresh  surface.    Care  !  2.   To  ascertain  whether  the  metal  is  fusible,  heat 
a  small  piece  of  it  in  a  stream  of  H.     3.  Is  it  heavier  or  lighter  than  water  ? 

(3)  i.  Expose  a  thin  slice  of  K  to  the  air.     What  takes  place?     2. 
Throw  a  small  piece  of  it  upon  H2O  (?).     In  this  experiment  it    is  ad- 
visable to  use  a  tall  beaker  and  to  cover  the  same  with  a  glass  plate.     3. 
What  is  the  action  of  the  halogens  upon  K?     Ask  for  instructions. 

POTASSIUM   AND    OXYGEN. 

(4)  Preparation  of  Potassium  Hydroxide. — In  an  iron  vessel  dissolve 
50  grms.  of  crystallized  Ba(OH)2  in  160  cc.  of  water.      Cautiously  add 
a  hot  concentrated  solution  of  20  grms.  of  K2SO4  until  a  sample  of  the 
supernatant  liquid  is  no  longer  precipitated  by  either  K2SO4  or  Ba(OH)2. 
Filter  rapidly  through  a  plaited  filter,  and  evaporate  the  solution  in  an 
iron  or  silver  dish  over  a  large  flame.     Continue  heating  the  residue  till  it 
appears  in  a  state  of  quiet  fusion.     During  this  operation  protect  the  eyes 
with  a  glass  plate.     Now  pour  the  product  upon  a  clean  iron  surface,  and 


36  EXPERIMENTS    IN    GENERAL   CHEMISTRY. 

while  still  warm  put  it  into  a  bottle  provided  with  a  well-fitting  stopper. 
Examine  its  fracture  and  color.  Try  its  solubility  in  water  and  in  alcohol. 
What  is  the  reaction  of  the  aqueous  solution  with  litmus  ?  What  is  an  alkali  ? 

Salts. 

(5)  Potassium  Chlorate.— KC1O3.     (See  p.  22). 

(6)  Potassium  Nitrate. — KNO3.     To  a  hot  concentrated  solution  of  20 
grms.  of  NaNO3  add  a  solution  of  18  grms.  of  KC1.     Boil.     What  sepa- 
rates from  the  warm  mixture  ?     What  crystallizes  from  the  mother  liquor 
on  cooling  ?  Recrystallize  the  latter  product.  Examine  its  crystalline  form. 
Is  it  more  soluble  in  hot  than  in  cold  water  ?     Explain  the  method   of 
preparation. 

(7)  Into  a  red-hot  platinum  crucible  throw  small  portions  of  an  intimate 
mixture  of  10  grms.  of  KNO3  and  i^  grms.  of  charcoal  powder.     What 
takes  place?     Write  the  reaction.     What  is  gunpowder  ? 

Reactions. 

(8)  Use  KNO3  for  the  following  tests. 

i.  Place  a  little  of  the  salt  upon  the  end  of  a  clean  platinum  wire  and 
introduce  it  into  a  non-luminous  flame.  What  color  do  you  observe? 
View  the  flame  through  a  cobalt  glass  (?).  2.  To  the  aqueous  solution 
of  the  potassium  salt  add  HCland  boil.  Concentrate  by  evaporation  and 
add  PtCl4.  What  is  the  composition  of  the  resulting  precipitate  ?  Try 
its  solubility  in  hot  and  in  cold  water,  also  in  alcohol.  3.  To  the  con- 
centrated solution  of  the  salt  add  a  saturated  solution  of  tartaric  acid  ; 
either  at  once,  or  on  shaking,  a  white  crystalline  precipitate  appears  (?). 

SODIUM.— Na. 

(1)  How  is  this  metal  usually  prepared  ? 

(2)  Study  its  physical  and   chemical  properties   (Richter,   p.    285). 
Wherein  does  it  differ  from  K? 

(3)  Prepare  Sodium  Amalgam. 

To  500  grms.  of  dry  mercury,  contained  in  a  Wedgewood  mortar  add 
gradually  5-10  grms.  of  Na  in  thin  slices.  Perform  this  operation  in  a 
good  draught  chamber,  as  the  union  of  the  two  metals  is  attended  with 
the  evolution  of  light  and  heat,  and  poisonous  vapors  are  given  off.  Stir 
well  with  the  pestle,  allow  to  cool,  and  transfer  the  product  to  a  well- 
stoppered  bottle.  What  is  its  action  on  H2O  or  dilute  H2SO4  ? 

SODIUM   AND    OXYGEN. 

(4)  Preparation  of  Sodium  Hydroxide  solution. 

Add  a  little  water  to  10  grms  of  fresh  quicklime  contained  in  an  iron 
(or  porcelain)  vessel.  Cover  the  latter,  and  in  a  second  iron  pot  dis- 
solve 25  grms.  of  soda  ash  (Na2CO3),  using  about  100  cc.  of  water. 


METALS    OF    THE   ALKALIES — SODIUM   AND    OXYGEN.  37 

Heat  the  solution  to  boiling ;  stir  the  quicklime — which  should  have 
broken  up  to  a  white  powder — with  enough  water  to  form  a  thin  paste 
(milk  of  lime),  and  add  this  gradually  to  the  boiling  liquid.  Stir  well 
with  an  iron  wire  ;  transfer  the  mixture  to  a  bottle ;  cork,  and  allow  it  to 
stand.  After  the  supernatant  liquid  has  become  perfectly  clear,  decant  it 
by  means  of  a  glass  siphon  filled  with  water.  It  should  be  preserved  in 
a  tightly  corked  bottle  (?).  Test  a  few  drops  of  the  solution  with  BaCl2  (?). 
What  should  the  solution  contain,  and  of  what  does  the  precipitate,  from 
which  it  was  separated,  consist?  Write  the  equation  representing  the  re- 
action. 

(5)  Determine  the  amount  of  NaOH  contained  in  the  solution. 
Measure  off  accurately  20  cc.  into  a  porcelain  dish ;  add  a  drop  or  two 

of  phenolphthalein  solution,  and  dilute  with  water.  From  a  burette  care- 
fully add  dilute  hydrochloric  acid  until  the  red  color  has  just  disappeared. 
Read  off  the  volume  of  the  acid  used ;  it  is  the  exact  quantity  needed  to 
neutralize  the  alkali : — 

NaOH  -f  HC1  =  NaCl  +  H2O; 

that  is,  40  pts.  (i  mol.)  of  NaOH  require  36.5  pts.  (i  mol.)  of  HC1, 
and  if  we  know  the  weight  of  the  HC1  contained  in  the  volume  of  the 
dilute  acid  consumed,  a  simple  proportion  will  give  the  weight  of  the 
alkali  in  20  cc.  of  the  solution.  The  strength  of  the  acid  is  determined 
as  follows:  In  a  porcelain  dish,  dissolve  1.06  grms.  of  pure  Na2CO3, 
previously  ignited  and  accurately  weighed;  add  a  little  phenolphthalein, 
heat  to  boiling  and  introduce  acid  from  the  burette  until  the  liquid 
remains  colorless  after  continued  boiling.  The  carbonate  is  then  exactly 
neutralized :  — 

Na2CO3  +  2HC1  =  2NaCl  +  CO2  +  H2O. 

It  takes,  therefore,  73  pts.  of  HC1  for  106  pts.  of  Na2CO3.  Suppose,  now, 
20  cc.  of  the  acid  had  been  used  to  decolorize  the  indicator,  then  i  cc. 
would  equal  iffi  =  .53  grms.  of  Na2CO3,  or  .365  grms.  of  HC1.  The 
latter  number  is  the  standard  or  strength  of  the  dilute  acid. 

The  phenolphthalein  takes  no  part  in  these  reactions ;  it  merely  indi- 
cates by  its  change  of  color  the  complete  neutralization  of  the  alkali. 
Why  is  it  necessary  to  boil  the  solution  when  the  acid  is  standardized 
with  a  carbonate  ? 

Salts. 

(6)  Sodium  chloride. — NaCl. 

Purify  common  salt. — Grind  50  grms.  of  salt  in  a  mortar  with  150  cc. 
of  water.  Filter  into  a  beaker,  and  conduct  HC1  gas  into  the  solution, 


38  EXPERIMENTS    IN    GENERAL   CHEMISTRY. 

as  shown  in  Fig.  36.      Pure  NaCl  separates  out.     Collect  it  on  a  platinum 
no.  36.  cone,  remove  the  liquid  with  the  aid  of  a  filter 

pump,  and  dry  the  salt  by  warming  it  in  a  porce- 
lain dish,  while  stirring  it  with  a  glass  rod. 
(7)  Sodium  carbonate. — Na2CO3. 
Recrystallize  some  of   the   commercial   carbo- 
nate.    Heat  a  portion  of  the  product  in  a  porce- 
lain dish.     What  do  you  observe  ? 

Reactions. 

(8)  Use  the  purified  chloride  for  the  tests,  i.  What  color  do  sodium 
salts  give  to  the  flame?  2.  Mix  a  drop  of  the  aqueous  solution  with  10 
drops  of  a  PtCl4  solution  on  a  watch-glass.  Evaporate  very  carefully  to 
a  small  volume.  On  cooling,  a  red  colored  salt  crystallizes  out  in  long 
monoclinic  needles  (?).  Is  it  soluble  in  water?  in  alcohol?  3.  Are 
there  any  salts  of  sodium  which  are  not  soluble  in  water  ?  Can  com- 
pounds of  sodium  be  precipitated  by  any  reagent  ? 

AMMONIUM. 

(1)  What  is  the  composition  of  ammonium?     Can  it  be  obtained  in  a 
free  state?     (See  Richter,  p.  295.) 

(2)  Dissolve  commercial  sal  ammoniac  in  a  little  water,  add  ammonia 
in  slight  excess,  warm,  filter  if  a  precipitate  is  formed,  and  evaporate  to 
crystallization;  stir  constantly.     Ammonium  chloride  is  thus  obtained 
in  the  form  of  a  fine  powder. 

Reactions. 

(3)  i.  On  a  piece  of  platinum  foil  heat  successively  small  portions  of 
the  chloride,  the  sulphate,  and  the  nitrate.     What  occurs  in  each  case  ? 
2.  Mix  a  little  NH4C1  with  burnt  lime  in  a  small  mortar.     Note  the  odor 
of  the  escaping  gas  and  its  reaction  with  litmus.     3.   Heat  a  small  por- 
tion of  NH4C1  with  a  caustic  soda  solution.     What  is  given  off?    Explain 
the  action  of  strong  bases  upon  ammonium  salts.     4.  Add  PtCl4  to  a 
solution  of   NH4C1.     Result?     5.  To   a   concentrated  solution   of   the 
ammonium  salt  add  tartaric  acid  and  shake  the  mixture  (?).     6.  Do  com- 
pounds of  ammonium  impart  a  color  to  the  flame  ? 


Compare  the  metals  of  the  alkalies  with  each  other.     How  can  the  com- 
pounds of  potassium,  sodium,  and  ammonium  be  distinguished  ? 

Problems. — i.  How  much  KNO3  is  theoretically  obtainable  from  2  kilos 


METALS    OF   THE    ALKALINE    EARTHS — CALCIUM.  39 

of  Chili  saltpetre  of  97%,  and  what  amount  of  Sylvite  containing  98%  of 
KC1  is  required?  2.  Suppose  that  75  cc.  of  dilute  HNO3  were  required 
to  saturate  50  cc.  of  a  potash  lye;  further,  that  10  cc.  of  the  acid  neu- 
tralized i. 06  grms.  of  Na2CO3,  what  amount  of  KOH  would  the  lye  con- 
tain? 3.  In  the  valuation  of  a  pearl  ash  (impure  K2CO3),  29.1  cc.  of  a 
sulphuric  acid  were  used  to  neutralize  5  grms.  of  the  sample  ;  the  acid 
contained  98  grms.  of  H2SO4  per  litre ;  calculate  the  percentage  of  im- 
purities in  the  product.  4.  Required  the  minimum  amount  of  marble  that 
should  be  burnt  to  liberate  the  NH3  from  50  grms.  of  NH4NO3. 


CHAPTER   VIII. 

METALS  OF  THE  ALKALINE  EARTHS— CALCIUM,  STRONTIUM,  BARIUM. 

CALCIUM.— Ca. 
CALCIUM   AND    OXYGEN. 

(1)  i.   Ignite  2  grms.  of  powdered  marble  in  a  platinum  crucible  to 
the  highest  temperature  obtainable  with  the  aid  of  the  blast  lamp.     Con- 
tinue this  for  15  minutes,  occasionally  stirring  the   mass  with  a  platinum 
wire ;  what  is  the  residue  ?     Explain  the  reaction.      2.  Add  about  5  cc. 
of  water  to  the  product.     What  do   you  observe?     Test  the  reaction  of 
the  product  with  litmus  paper. 

(2)  i.   Prepare  lime  water. — To    the   slaked    lime  obtained    from  20 
grms.  of  quicklime  (see  p.  37)  add  i  litre  of  water ;    transfer  the  mixture 
to  a  bottle.     Cork  tightly,  shake  and  allow  to  stand.     When  the  solution 
has  become  clear,  draw  it  off  by  means  of  a  siphon  ?     What  does  it  con- 
tain ?     Of  what  does  the  undissolved  portion  consist  ?     2.   Place  a  por- 
tion of  the  lime  water  on  a  watch  glass  and  expose  to  the  air(?).     3. 
Through  a  second  portion  blow  air  from  your  lungs  (?).     4.   Conduct  a 
stream  of  CO2  through  a  third  portion  and  observe  carefully  the  successive 
changes.     Explain  them.    5.  What  takes  place  upon  boiling  the  clear  solu- 
tion which  is  obtained  as  the  final  product  in  the  preceding  experiment  ? 

Salts. 

(3)  Calcium  Chloride.— CaCl2. 

i.  Evaporate  some  of  the  spent  acid  of  a  CO2  generator  to  dryness. 
What  is  the  residue?  2.  Expose  a  little  of  the  salt  to  the  air  (?).  3. 
What  use  have  you  made  of  CaCl2  previously?  4.  Prepare  porous  CaC/^ 
(CaCl2  -f-  2H2O).  Dissolve  the  residue  obtained  in  i  in  lime  water, 
filter,  and  neutralize  exactly  with  HC1.  Evaporate  the  filtrate  to  dryness 


40  EXPERIMENTS    IN    GENERAL    CHEMISTRY. 

in  a  porcelain  dish,  and  heat  the  residue  for  some  time  on  the  sand-bath. 
The  solution  of  the  product  must  show  a  neutral  reaction. 

(4)  Calcium  Hypochlorite.—C*(C\Q\.     (See  p.  22.) 

(5)  Calcium  Sulphate.— CaSO4. 

i.  Carefully  heat  a  few  grms.  of  gypsum  in  a  porcelain  dish  until  the 
water  of  crystallization  is  completely  expelled.  Pulverize  the  residue. 
What  happens  when  it  is  made  into  a  paste  with  water  and  allowed  to 
stand  ? 

Reactions. 

Use  the  pure  CaCl2  for  the  following  tests : — 

i.  Introduce  a  small  portion  of  the  salt  into  the  Bunsen  flame  by 
means  of  a  platinum  wire  (?).  2.  To  the  aqueous  solution  add  (NH4)2 
CO3.  Result  ?  3.  To  another  portion  add  dilute  H2SO4.  What  is  the 
composition  of  the  precipitate?  Why  does  it  not  form  in  very  dilute 
solutions?  3.  Add  (NH4)2C2O4  to  the  filtrate  from  the  CaSO4  (?). 
*  t 

STRONTIUM.— Sr. 

Reactions. 

i.  What  color  is  imparted  to  the  Bunsen  flame  by  compounds  of  this 
element  ?  2.  Add  a  CaSO4  solution  to  the  solution  of  a  strontium  salt  (?). 

BARIUM.— Ba. 

Reactions. 

i.  Observe  what  color  Ba  compounds  give  to  the  flame.  Moisten  the 
sample  with  HC1  before  heating  it  (?).  2.  To  a  portion  of  the  aqueous 
solution  of  the  chloride  add  (NH4)2CO3.  What  results?  2.  Add  dilute 
H2SO4  to  a  second  portion  (?). 


Point  out  how  the  elements  of  this  group  may  be  distinguished  (#) 
from  those  of  the  preceding  group ;  (b}  from  each  other. 

Problems. — i.  How  much  nitric  acid  of  20  per  cent,  will  effect  the  so- 
lution of  i  grm.  of  Iceland  spar  (CaCO3)  ?  How  much  CO2  is  given  off, 
and  what  volume  would  it  occupy  at  20°  C.  under  a  pressure  of  750  mm.  ? 
2.  Suppose  .5  grm.  of  sulphur  were  dissolved  in  HNO3,  what  quantity 
of  BaCl2  must  be  added  until  it  ceases  to  produce  a  precipitate  ?  3.  One 
grm.  of  a  mineral  consisting  of  the  carbonates  of  Ca,  Sr,  and  Ba,  in  the 
proportion  of  their  molecular  weights,  will  leave  what  weight  of  the  mixed 
sulphates  on  treating  and  evaporating  with  an  excess  of  H2SO4  ? 


MAGNESIUM    GROUP MAGNESIUM.  41 

CHAPTER  IX. 

MAGNESIUM    GROUP— MAGNESIUM,    ZINC,   CADMIUM. 
MAGNESIUM.— Mg. 

(1)  Examine  the   metal  in  the  forms  of  ingot,  ribbon  and   powder* 
Note  its  color,  lustre  and  specific  gravity.     2.  Introduce  a  piece  of  the  rib- 
bon into  the  flame  with  the  forceps  (?).     What  is  the  product  ?    3.   Treat 
a  piece  of  the  ribbon  with  dilute  H2SO4.     Reaction  ? 

Salts. 

(2)  Magnesium  Chloride. — MgCl2. 

Prepare  the  ANHYDROUS  salt. — Dissolve  about  50  grms.  of  the  crystal- 
lized (?)  chloride  and  50  grms.  of  NH4C1  in  as  little  water  as  possible. 
Evaporate  to  dryness  in  a  porcelain  dish.  Reduce  the  mass  while  hot  to 
small  pieces  in  a  mortar,  dry  it  carefully,  so  as  to  remove  every  trace  of 
moisture.  It  is  best  to  do  this  by  heating  small  portions  of  the  material 
in  a  porcelain  crucible  until  it  no  longer  sinters.  A  small  sample  should 
not  give  off  moisture  when  heated  in  a  dry  test-tube.  Be  careful  also  to 
prevent  re-absorption  of  moisture.  Quickly  transfer  the  warm  powder  to 
a  platinum  crucible  provided  with  a  well-fitting  cover.  Heat,  at  first 
gently,  to  expel  the  NH4Cl,  then  increase  the  temperature  until  the  mass 
is  in  a  state  of  quiet  fusion.  It  is  the  anhydrous  salt  which,  being 
extremely  hygroscopic,  should  be  preserved  in  a  tightly  stoppered  bottle. 
It  should  dissolve  in  water  to  a  clear  liquid. 

Why  cannot  the  anhydrous  chloride  be  obtained  by  evaporation  of  the 
aqueous  solution  ? 

(3)  Magnesium  Sulphate. — Mg  SO4  -f  yH2O. 

Recrystallize  some  of  the  commercial  salt.  What  is  the  form  of  the 
crystals  ?  Taste  ? 

(4)  Reactions. 

i.  Heat  a  portion  of  the  sulphate  or  chloride  on  a  platinum  wire  in 
the  Bunsen  flame ;  moisten  with  Co(NO3)2  solution  and  heat  again.  A 
pink-colored  mass  results.  2.  Add  some  caustic  soda  to  a  little  of  the 
solution  of  the  chloride  (?).  The  resulting  precipitate  dissolves  on  ad- 
dition of  an  ammonium  salt  (?)  3.  Mix  a  second  portion  of  the  chloride 
solution  with  NH3  and  NH4C1,  add  Na2HPO4  and  agitate  the  liquid.  What 
is  the  composition  of  the  precipitate  ?  Examine  it  with  the  aid  of  a  lens. 
6 


42  EXPERIMENTS    IN    GENERAL   CHEMISTRY. 

ZINC.— Zn. 

(1)  How  is  this  metal  obtained  from  its  ores? 

(2)  Study  the   physical  and  chemical  properties  of  Zn  (see  Richter,  p. 
316).      i.  Treat  a  small  piece  of  pure  Zn  with  dilute  H2SO4(?).      2.   Re- 
peat this  experiment,  substituting  the  impure  commercial  metal.     What 
difference  do  you  observe  ?     What  causes  it  ? 

(3)  Granulate  commercial  zinc. — Melt  100  grms.  of  the  metal  in  a  well- 
covered  Hessian  crucible.     The  blast  lamp  maybe  used  for  this  purpose, 
but  it  is  better  to  perform  the  operation  in  a  wind  furnace.     The  crucible 
is  then  removed  from  the  source  of  heat,  and  allowed  to   cool  until  the 
melted  metal  no  longer  takes  fire  when   the  cover  is  lifted.     Pour  the 
metal,  in   a  thin   stream,  into  a  pail  filled  with  cold  water.     Drain  the 
product  and  dry  at  a  moderate  heat. 

Salts. 

(4)  Zinc  sulphate. — Zn  SO4  -f  7H2O.  (See  p.  12).      i.   Prepare  some 
of  this  salt  and  recrystallize    it   carefully  from    water.      2.   Examine  the 
crystals.     What  other  salt  have  you  prepared  that  exhibits  similar  forms  ? 
Is  there  any  analogy  in  the  composition  of  the  two  salts  ? 

Reactions. 

(5)  i.   Heat  a  small  piece  of  Zn  on  charcoal  in  the  oxidizing  flame. (?) 
2.  Moisten  the  incrustation  obtained  with  a  drop  of  Co(NO3)2,  and  heat 
again.     Result  ?     3.  To  a  solution  of  ZnSO4  add  (NH4)2  S.     What  is  the 
color  of  the  precipitate  ?  Try  its  solubility  in  dilute  HC1  and  in  HC2H3O, 
(acetic  acid).     4.  Study  the  action  of  caustic  alkalies,  e.  g.,  NaOH  upon 
the  Zn  solution. 


How  could  you  distinguish  between  Zn  and  Mg  ?  What  differences 
are  there  between  this  and  the  preceding  groups  ? 

Problems. — i.  What  is  the  strength  of  a  sulphuric  acid  of  which  20  cc. 
dissolve  exactly  .048  grm.  of  Mg?  2.  Suppose  it  was  found  that  i  grm. 
of  Zn  gave  with  H2SO4,  325  cc.  of  H  at  16°  C.  and  755  mm.,  and, 
further,  that  .369  grm.  of  Mg  produced  the  same  amount  of  the  gas. 
Knowing  the  atomic  weight  of  Mg  to  be  24,  and  remembering  that  the 
two  sulphates  are  isomorphous,  how  is  it  possible  to  deduce  the  at.  wt.  of 
Zn  from  the  data  given  ? 


MERCURY   AND    OXYGEN.  43 

CHAPTER  X. 

MERCURY,  COPPER,  SILVER,  GOLD. 
MERCURY.— Hg. 

(i)  Study  the  physical  and  chemical  properties  of  the  metal.  Wherein 
does  it  differ  from  the  other  metals? 

MERCURY   AND    OXYGEN. 

(  2 )  Mercuric  oxide.  — Hg  O . 

How  is  this  substance  prepared?     What  is  its  behavior  on  heating? 

Mix  a  little  powdered  S  with  dry  Na.2CO3  and  HgO.  Ignite  the  mix- 
ture in  a  dry  test-tube.  Extract  the  residue  with  water,  filter,  acidify 
with  HC1  and  add  BaCl2.  What  has  become  of  the  oxide  of  mercury  in 
this  experiment? 

Salts. 

(3)  Mercurous  Nitrate. — HgNO3. 

An  excess  of  metallic  mercury  (use  10-15  grmsO  ig  treated  in  the  cold 
with  moderately  strong  HNO3  until  the  formation  of  crystals  is  no  longer 
noticeable.  Redissolve  the  crystals  by  warming,  filter,  and  allow  to 
crystallize. 

To  prepare  a  solution  of  the  salt  take  it  up  with  water  acidulated  with 
HN03  (?). 

(4)  Mercuric  chloride. — HgCl2. 

Dissolve  about  5  grms.  of  Hg  in  aqua  regia.  Evaporate  to  dryness  on 
a  water  bath.  Place  the  residue  into  a  small  dry  flask,  cover  the  latter 
with  a  watch-glass,  and  heat  cautiously  on  a  sand-bath.  What  is  the 
sublimate  formed  in  the  upper  part  of  the  flask  ?  Dissolve  it  in  four  parts 
of  boiling  water  and  allow  to  crystallize. 

Reactions. 

(5)  Mercurous  compounds.     Use  the  solution  of  the  nitrate,      i.  Add 
a  few  drops  of  HC1  to  2  or  3  cc.  of  the  solution.     What  takes  place  ? 
Filter,  and  add  NH3  to  the  precipitate  (?).     2.   Add  stannous  chloride 
to  another  portion  of  the  nitrate  solution  (?).     3.  In   a  third  portion 
immerse  a  slip  of  Cu  foil.     Examine  the  stain  on  the  metal ;  is  it  changed 
when  you  hold  it  in  the  flame? 

(6)  Mercuric  compounds.     The  chloride  will  answer  for  the  tests. 

i.  Pass  H2S  through  a  dilute  solution  and  observe  the  gradation  of 
colors  through  which  the  precipitate  passes.  What  is  the  final  product  ? 


44  EXPERIMENTS    IN    GENERAL   CHEMISTRY. 

2.  Add  SnCl2,  drop  by  drop,  to  the  mercury  solution.  Explain  the 
changes  which  occur. 

COPPER.— Cu. 

i.  Preparation. — Ignite  the  pure  oxide  in  a  current  of  dry  H  (see  p. 
14).  Examine  the  color  and  the  lustre  of  the  product  \  test  its  solu- 
bility in  HC1,  H2SO4  (both  strong  and  dilute),  and  HNO3.  Write  equa- 
tions representing  the  reactions. 

Salts. 

(2)  Copper  Sulphate.— CuSO,  -f  sH2O. 

To  10  grms.  of  Cu  in  a  flask  add  45  grms.  of  cone.  H2SO4,  and  heat. 
When  the  metal  has  completely  disappeared  and  the  gas  (?)  ceases  to  be 
given  off,  allow  to  cool,  place  the  white  crystalline  residue  (?)  into  a 
porcelain  dish,  rinse  the  flask  with  hot  water.  Now  add  a  few  drops  of 
HNO3  to  the  hot  water  solution,  and  filter.  From  the  filtrate  the  sulphate 
crystallizes  on  standing.  Recrystallize  the  product. 

Does  this  salt  suffer  decomposition  on  exposure  to  the  atmosphere? 
Heat  a  small  quantity  in  a  porcelain  crucible,  first  moderately,  then  more 
strongly  (?). 

(3)  Sulphate  of  Copper  and  Potassium.— CuK2(SOt)2  -f-  6  H2O. 
Prepare  solutions  of  10  grms.  of  blue  vitriol  and   7  grms.  of  K2SO4, 

both  saturated  at  70°.  The  latter  should  also  contain  a  few  drops  of 
H2SO4.  Mix  the  solutions ;  on  cooling  the  double  salt  separates  in 
whitish-blue  crystals.  Examine  their  form. 

Reactions. 

(4)  Use  either  of  the  salts  you  have  prepared. 

i.  Mix  a  little  of  the  salt  with  Na2CO3,  and  heat  on  charcoal  in  the 
reducing  flame  (?).  2.  Make  a  borax  bead  and  dissolve  a  minute  quan- 
tity of  a  Cu-compound  in  it.  What  color  does  it  give  (a)  in  the  oxidiz- 
ing flame?  (<£)  in  the  reducing  flame?  (r)  when  the  bead  is  reduced  with 
a  small  piece  of  tin  ?  3.  Through  a  dilute  Cu-solution  pass  H2S.  Is  the 
resulting  precipitate  soluble  in  HC1  or  in  HNO3?  4.  Add  ammonia, 
drop  by  drop,  to  the  solution.  What  changes  do  you  observe?  5.  To 
a  portion  of  the  very  dilute  solution  add  potassium  ferrocyanide  (?). 

(5)  To  a  solution  of  copper  sulphate  in  a  porcelain  dish  add  a  small 
piece  of  Zn.     Allow  to  stand  over  night.     Note  the  result.     Has  the  Zn 
disappeared  ?     Does  the  solution  contain  any  of  this  metal  ?     In  what 
form  ?     Where  is  the  Cu  ? 

(6)  Repeat  the  experiment,  weighing  the  copper  sulphate  (.5  grm.) 
and  the  Zn  (.2  grm.).     Add.HCl  in  quantity  sufficient  to    insure  the 


SILVER.  45 

entire  solution  of  the  Zn,  collect  the  Cu  on  a  filter,  wash  with  alcohol, 
dry,  heat  gently  and  weigh  it  in  a  porcelain  crucible.  The  filtrate  should 
be  colorless. 

Compare  the  weight  of  the  metallic  Cu  obtained  with  that  of  the  Zn 
employed  (?).  How  does  the  found  Cu  accord  with  the  calculated  amount 
of  that  metal  in  .5  grm.  CuSO4.5H2O? 

Repeat  the  experiment  using  Cd  in  place  of  Zn.  Compare  the  weights 
of  the  metals  as  before.  What  deduction  can  you  make  ? 

SILVER.— Ag. 

(1)  Prepare  pure  Silver  from  a  coin. 

Dissolve  a  25-cent  piece  in  nitric  acid  of  sp.  gr.  1.2,  filter  (?),  and 
evaporate  the  blue  (?)  solution  to  dryness.  Fuse  the  residue  till  it 
blackens,  extract  with  250  cc.  of  water ;  filter.  Now  add  ammonia  in 
large  excess,  and  then,  cautiously,  a  sodium  bisulphite  solution  (of  about 
40  %)  until  on  boiling  a  small  portion  of  the  liquid,  it  is  completely 
decolorized. 

The  greater  part  of  the  Ag  separates  from  the  solution  on  standing  in 
the  cold  ;  it  is  well  crystallized.  The  remainder  may  be  precipitated  by 
warming  to  70°.  Digest  the  product  with  strong  ammonia  (?),  wash,  dry 
and  ignite  it. 

Examine  the  metal  carefully.  What  are  its  physical  and  chemical 
characteristics? 

SILVER   AND    SULPHUR. 

(2)  Silver  Sulphide.—  Ag.2S. 

Into  a  dilute  solution  of  AgNO3  (see  next  experiment),  containing 
about  2  grms.  of  the  metal,  pass  H.2S.  When  the  liquid  smells  of  the  gas, 
filter  off  the  black  precipitate,  wash  it  with  water  and  dry  at  100°. 

Salts. 

(3)  Silver  Nitrate. — AgNO3. 

Dissolve  the  Ag  obtained  in  (i)  in  dilute  HNOS  and  evaporate  to  dry- 
ness  on  the  water  bath.  Dissolve  the  residue  in  80  cc.  of  distilled  water, 
and  preserve  the  solution  in  a  dark  bottle  (?).  What  is  its  reaction  with 
litmus  ? 

Reactions. 

(4)  i.    Compounds  of  Ag  on    charcoal  before    the  blow-pipe  give  a 
white  metallic  globule  (?).     2.  To  a  silver  solution — use  the  nitrate — add 
HC1.     Collect  the  precipitate  on  a  small  filter,  wash,  dissolve  it  in  am- 
monia, and  add  an  excess  of  HNO3  to  the  solution  (?).     Explain  these 


46  EXPERIMENTS    IN    GENERAL   CHEMISTRY. 

reactions.  3.  Expose  a  small  portion  of  the  chloride  to  direct  sunlight. 
Any  change  ?  What  practical  application  is  made  of  this  reaction  ? 
(Read  Richter,  p.  340.)* 

(5)  Place  strips  of  the  metals  Zn,  Fe,  Sn,  Pb,  and  Cd  in  a  solution  of 
silver  nitrate.     What  is  the  result  in  each  case?     Explain. 

GOLD.— Au. 

(1)  How  could  you  distinguish  the  metal  Au  from  Hg,  Ag,  and  Cu? 

Reactions. 

(2)  i.  Dissolve  a  small  piece  of  gold  (or  of  a  substance  containing 
gold)  in  aqua  regia,  concentrate  the  solution  at  a  gentle  heat  and  pour  it 
into  a  porcelain  dish.    Add  a  solution  of  FeCl3  to  an  SnCl2  solution  until 
the  latter  is  permanently  yellow.     After  diluting,  dip  a  glass  rod  into  this 
and  then  into  the  gold  solution.     A  purple  streak  (purple  of  Cassius)  is 
formed.      2.  Add  ferrous  sulphate  to  some  of  the  AuCl3  solution  (?). 


In  what  respects  do  the  members  of  this  group  differ  from  each  other, 
and  how  can  they  be  distinguished  from  the  metals  of  the  preceding 
groups  ? 

Problems. — 1.5  grms.  of  HgO  gave  on  ignition  with  carbon  4.63  grms. 
of  metallic  mercury;  the  specific  gravity  of  the  vapor  of  HgCl2  referred 
to  H,  was  found  to  be  135.5.  What  is  the  atomic  weight  of  Hg  ?  2. 
The  molecule  of  Hg  contains  how  many  atoms,  if  the  vapor  density 
equals-  TOO?  3.  On  analysis  a  chalcocite  was  found  to  contain  20.15  per 
cent,  of  S  and  79.85  per  cent,  of  Cu.  Deduce  the  molecular  formula 
of  the  mineral.  4.  What  quantities  of  Ag,  Au,  and  Hg  can  be  precipi- 
tated from  their  respective  solutions  by  i  grm.  of  Cu  ? 


CHAPTER   XI. 

ALUMINIUM,  TIN,  LEAD,  BISMUTH. 

ALUMINIUM.— Al. 

(i)  By  what  methods  is  this  metal  obtained  on  a  large  scale?  What 
are  its  properties?  Try  the  action  of  the  following  reagents  upon  Al : 
HC1,  HNO3,  and  NaOH  solution.  Write  the  reactions. 

*  If  practicable,  the  instructor  should  here  show  and  explain  the  preparation  of  a 
photographic  negative. 


TIN.  47 

Sa/ts. 

(2)  Sulphate  of  Aluminium  and  Potassium. — KA1(SO4)2  -j-  i2H2O. 

Prepare  saturated  solutions  of  A12(SO4)3  and  K2SO4 ;  mix  these  so  that 
the  resulting  liquid  contains  the  two  sulphates  approximately  in  the  pro- 
portion of  their  molecular  weights.  The  double  sulphate  crystallizes  on 
standing.  Why?  Recrystallize  it  from  water.  What  is  the  form  of  the 
crystals  ? 

What  is  an  alum!     (See  Richter,  p.  351.) 


Reactions. 

(3)  Use  alum.  i.  Heat  a  little  of  the  salt  on  a  platinum  wire  in  the 
oxidizing  flame,  moisten  with  Co(NO3)2,  and  heat  again.  A  blue  mass 
(?)  is  the  product.  2.  To  an  aqueous  solution  add  ammonia  (?).  Add 
(NH4)2S  to  another  portion  of  the  solution.  What  do  you  observe  ?  4. 
To  the  diluted  solution  add  NaOH,  drop  by  drop.  Note  the  successive 
changes  (?). 


TIN.— Sn. 

(1)  Examine   a  bar  of  this  metal,      i.  Note  the  sound   it    emits   on 
bending  (?).      2.   Etch  a  smooth  surface  with  HC1  (?).     3.   Try  the  solu- 
bility of  Sn  in  hot  HC1.     4.   What  action  have  moderately  dilute,  and 
concentrated,  HNO3  upon  it?     Write  the  reactions. 

(2)  Determine  the  specific  heat  of  Tin. 

A  thin  glass  beaker  of  about  200  cc.  capacity  is  carefully  covered  on 
the  outside  with  a  moderately  thin  layer  of  cotton  wool.  This  may' be 
called  the  calorimeter.  Pour  100  cc.  of  distilled  water  into  the  beaker. 
Suspend  a  thermometer  in  the  water.  Place  25  grms.  of  granulated  tin 
into  a  test-tube,  close  the  mouth  of  the  latter  with  a  plug  of  cotton. 
Introduce  the  test-tube  with  its  contents  into  a  beaker  glass  containing 
boiling  water.  A  stout  copper  wire  will  serve  as  a  handle.  After  ten  or 
fifteen  minutes  the  tin  will  have  acquired  the  temperature  of  the  boiling 
water — 100°.  The  tube  is  then  rapidly  removed  from  the  latter  and  its 
outer  surface  freed  from  moisture  by  quickly  passing  a  towel  over  it. 
Remove  the  cotton  from  the  mouth,  and  transfer  the  tin  to  the  calori- 
meter. While  the  metal  is  being  introduced  raise  the  thermometer  from 
the  water,  and  replace  it  as  soon  as  all  the  metal  has  been  added ;  stir  the 
liquid  well  and  observe,  as  accurately  as  possible,  the  highest  point 


48  EXPERIMENTS    IN    GENERAL    CHEMISTRY. 

reached  by  the  mercury  column.     Approximate  results  can  be  obtained 
from  these  data.     Calculate  as  follows  :— 

Let  y  —  temperature  of  water  before  introducing  the  tin. 
"    z  —  "  "       "      after  " 

"  w  =  weight  of  the  water. 
"  v  =  "  of  the  metal. 
"  x  =  sp.  heat — then 

x__  loo(z-y) 
"25  (loo-z)   • 

(Study  Richter  pp.  256-259.)  Would  the  specific  heat  found  for  tin, 
when  multiplied  by  the  constant  6.4  give  the  same  value  as  that  found  in 
experiment  (3)  for  the  equivalent  of  tin  ?  Explain.  How  many  series 
of  tin  compounds  are  there  ? 

(3)  Determine  the  equivalent  weight  of  Tin. 

Place  about  3  grms.  of  tin  in  a  porcelain  crucible  that  has  been  pre- 
viously weighed.  Cover  the  metal  with  5-10  cc.  of  concentrated  HNO3. 
Then  carefully  apply  heat  by  means  of  an  iron  plate.  The  tin  is  dis- 
solved, while  fumes  of  NO2  are  set  free.  When  the  acid  has  been  entirely 
expelled,  heat  the  crucible  with  the  white  stannic  oxide  over  a  Bunsen 
burner ;  allow  to  cool  and  weigh. 

Let  w  =  weight  of  crucible  and  SnO2 
"    v  =       "        "         "         "  metallic  tin. 
a.    v  _        ((        «         (t 
y  — 

Then  w  -  v  =  weight  of  O, 
andv  -  y  =       "       "  Sn. 

Equiv.  ofSn=i^-^8 

W—  V 

Softs. 

(4)  Stannous  Chloride. — SnCl2. 

Dissolve  10  grms.  of  granulated  Sn  in  warm  cone.  HC1  with  the  addition 
of  a  few  drops  of  PtCl4  (?).  Put  the  solution  into  a  well-stoppered  bottle. 

Reactions. 

(5)  Stannous  Compounds.     Use  the  chloride  solution. 

i.  Conduct  H2S  through  a  portion  of  the  diluted  liquid.  A  brown 
precipitate  (?)  is  thrown  down.  Is  it  soluble  in  yellow  ammonium  sul- 
phide? What  does  HC1  precipitate  from  the  sulphide  solution?  2. 
What  is  the  action  of  HgCl2  upon  SnCl2  (see  p.  44). 

(6)  Stannic  Compounds.     Add  a  few  drops  of  Br  to  a  portion  of  the 
SnCl2  solution,  and  boil  (?).     Use  the  diluted  liquid  for  the  tests. 

1.  Pass  H2S  into  a  portion  of  the  solution.     What  is  the  color  of  the 
precipitate?     Is  it  soluble  in  HC1?  in  (NH4)2S? 

2.  Add  Cu-turnings,  boil,  decant  the  liquid,  and  add  HgCl2.     What 
happens?     Explain.. 


BISMUTH.  49 


LEAD.— Pb. 

(i)  How  can  this  metal  be  obtained  from  the  oxide?  By  what  physi- 
cal properties  can  it  be  distinguished  from  other  metals  ?  Is  it  soluble  in 
the  mineral  acids?  (2)  In  a  solution  of  5  grms.  of  lead  nitrate  in 
about  50  cc.  of  water,  suspend  a  strip  of  metallic  Zn  and  let  stand  for  a 
few  days  (?). 

Salts. 

(3)  Dissolve  5  grms.    of  granulated  lead  (test-lead)  by  warming  with 
dilute  HNO3.     Concentrate  by  evaporation  and  allow  to  crystallize. 

Reactions. 

(4)  i.   Before  the  blowpipe,  on  charcoal,  lead  compounds  are  reduced 
to  metallic  beads,  which  are  sectile  with  the  knife.      2.  Add  HC1  to  a 
solution  of  the  nitrate.     Boil  the  precipitate  with  water.  (?)     What  takes 
place  on   cooling?     3.  To   another  portion  add    dilute  H2SO4  (?).     4. 
Pass  H2S  into  a  third  portion  (?). 

BISMUTH.— Bi. 
Reactions. 

(i)  i.  Mix  a  little  of  the  oxide  or  nitrate  of  Bi  with  Na2CO3  and  heat 
in  the  reducing  flame  on  charcoal.  Does  the  resulting  metallic  globule 
resemble  lead  ?  Is  it  sectile?  2.  Pass  H2S  into  a  solution  of  the  chloride 
or  nitrate  in  HC1  (?).  3.  Add  a  large  volume  of  water  to  a  bismuth 
solution.  What  occurs?  What  reactions  distinguish  Al,  Sn,  Pb  and  Bi 
from  each  other,  and  from  the  metals  previously  studied  ? 

Problems. — i.   What  is  the  molecular  formula  of  a  mineral  containing 

SiO2  =  43.08 
A1203  =  36.82 
CaO  =  20.10 

100.00 

2.  A  compound  of  tin  and  chlorine  yielded  on  analysis  29.42  parts  of 
Sn  and  35.40  parts  of  Cl ;  its  vapor  density  was  ascertained  to  be  132.85. 
What  is  the  atomic  weight  of  tin  ?  3.  Deduce  the  formula  of  Cosalite 
from  the  following  analysis  : — 

S  =  15.27 

Bi  =  41.76 

Pb  =  40.32 

Ag  =  2.65 


IOO.OO 


50  EXPERIMENTS    IN    GENERAL    CHEMISTRY. 

CHAPTER  XII. 

CHROMIUM,    MANGANESE,    IRON,    NICKEL/  COBALT. 
CHROMIUM.— Cr. 

CHROMIUM  AND  OXYGEN. 

(1)  Chromic  oxide. — Cr.2O3. 

i.  Preparation. — Mix  intimately  20  grms.  of  potassium  dichromate 
and  4  grms.  of  sulphur.  Heat  the  mixture  in  a  porcelain  crucible  over 
the  blast  lamp  for  about  20  minutes.  Cool,  extract  the  residue  with 
boiling  water  and  dry  it  at  a  gentle  heat.  What  is  its  color  ;  is  it  soluble 
in  dilute  HC1?  2.  Fuse  a  portion  of  it  with  six  times  its  weight  of 
NaHSO4  in  a  platinum  crucible.  What  takes  place?  3.  Repeat  this 
experiment  with  some  finely  powdered  chromite.  (?) 

Salts. 

(2)  Chromic  Chloride.— CrCl3. 

Prepare  the  anhydrous  salt.  Intimately  mix  10  grams  of  Cr2O3,  pre- 
pared as  described,  and  3  grms.  of  powdered  charcoal,  and  convert  this 
into  a  dough  with  a  little  starch  paste.  Form  the  product  into  balls  of  the 
size  of  a  pea  ;  dry,  and  then  ignite  these  (covered  with  charcoal  powder) 
in  a  Hessian  crucible,  provided  with  well-fitting  lid.  Place  the  residue 
into  a  tube  of  hard  glass,  and  heat  it  in  a  current  of  CO.2  to  expel  every 
trace  of  moisture.  With  the  aid  of  a  blast  lamp  increase  the  tempera- 
ture and  replace  the  CO2  by  a  current  of  Cl.  The  excess  of  Cl  should 
be  absorbed  by  conducting  it  into  a  bottle  filled  with  caustic  soda.  (?) 
The  resulting  CrCl3  sublimes  to  the  cooler  portions  of  the  tube.  Describe 
its  appearance.  Is  it  soluble  in  water? 

What  other  chlorides  are  prepared  in  a  similar  way  ?  Write  the  equa- 
tion, expressing  the  reaction. 

(3)  Chrome  Alum.— Cr2(SO4)3.K2SO4  +  24  H2O. 

Dissolve  10  grms.  of  K2Cr2O7  in  a  little  water  ;  acidify  with  H2SO4,  pass 
SO2  into  the  liquid  until  the  latter  is  saturated  with  the  gas.  Allow  to 
stand  ;  the  double  salt  crystallizes.  What  is  its  crystalline  form  ?  Dis- 
solve some  of  it  in  cold  water  and  note  the  color  of  the  solution  ;  now 
warm  it.  What  takes  place  (see  Richter,  p.  374)? 

(4)  i.  Examine  crystals  of  potassium  dichromate,  K2Cr2O7.     How  is  it 
obtained?     2.  Dissolve  10  grms.  of  this  salt  in  water,  and  from  a  burette 
carefully  add  a  caustic  soda  solution  until  the  color  is  changed  to  yellow  (?). 


MANGANESE   AND    OXYGEN.  51 

What  crystallizes  from  the  solution  on  evaporation  ?     How  can   you  re- 
convert the  product  into  the  dichromate  ? 

Reactions. 

(5)  i.   Dissolve  a  minute  quantity  of  a  chromium  compound  in  a  borax 
bead.     Heat    in    the    oxidizing  and   in   the  reducing  flame.     Results? 
2.   Heat  a  little  of  the  compound  with  KNO3  on  a  platinum  foil  (?) 

(6)  Chromic  compounds.     Use  chrome  alum   for  the  tests.      i.   Add 
caustic  soda,  drop  by  drop,  to  a  little  of  the  solution.  (?)     Continue  the 
addition  of  the  reagent  till  the  precipitate  is  redissolved.     What    takes 
place  on  boiling  the  solution  ?     2.  What  is  the  action  of  ammonia  on  the 
solution  of  the  chromium  salt  ? 

(7)  Chromates.     Use  a  solution  of  potassium  chromate.      i.   Add  lead 
acetate  solution.     Note   the  color    of  the  precipitate.     Is  it   soluble  in 
acetic  acid?     2.   Substitute  BaCl2  for  the  lead   salt  in  the   preceding  ex- 
periment. (?)     3.  Acidify  the  chromate  solution  with   H2SO4  and  add 
H2O2  to    the    liquid.     What   happens?      4.   To  some  of  the   chromate 
solution  add  a  few  drops  of  HC1  and  about  i  cc.  of  alcohol.     What  occurs 
when  the  mixture  is  heated  to  boiling  ? 

MANGANESE.— Mn. 
MANGANESE    AND    OXYGEN. 

(1)  In  what  proportions  do  these  two  elements  unite  with  each   other? 
Enumerate  the  oxides  which  occur  in  nature.     What  is  formed  when  the 
oxides  of  manganese  are  heated  in  H  ?     When   they  are  ignited   in  the 
air? 

Salts. 

(2)  Manganous  Chloride. — MnCl2  -f-  4H2O. 

Evaporate  in  a  porcelain  dish  the  solution  obtained  in  the  preparation 
of  Cl  from  MnO2  and  HC1.  Heat  the  dry  residue  over  a  small  flame  for 
some  time.  Add  much  water  and  boil.  Filter,  and  to  ^  of  the  filtrate 
add  a  solution  of  Na2CO3  in  excess.  Allow  the  precipitate  (?)  to  settle, 
draw  off  the  supernatant  liquid  with  a  siphon,  and  wash  the  remaining 
precipitate  several  times  with  water  by  decantation.  Add  the  precipitate 
then  to  the  principal  solution  and  digest  at  a  gentle  heat  until  a  small 
filtered  sample  mixed  with  (NH4)2S  gives  a  flesh-colored  precipitate  which 
is  completely  dissolved  by  dilute  acetic  acid.  Now  filter  and  evaporate 
to  crystallization. 

(3)  Potassium  Manganate — K2MnO4  and 

Potassium  Permanganate — K2Mn2O8. 


52  EXPERIMENTS    IN    GENERAL   CHEMISTRY. 

In  a  porcelain  crucible  fuse  a  mixture  of  5  grms.  KOH  and  2.5  grms. 
KC1O3  \  gradually  add  5  grms.  finely  powdered  MnO2.  Maintain  a 
moderate  red  heat  for  15  minutes.  Dissolve  the  dark-green  residue  in  a 
little  water.  Observe  the  color  of  the  solution.  What  does  it  contain? 
Then  dilute  with  much  water  and  conduct  CO2  into  the  liquid.  Is  there 
any  change  ?  If  so,  write  the  equation  expressing  it. 

K2Mn2O8  as  well  as  K2MnO4  are  powerful  oxidizing  agents.  Warm  a 
little  of  the  alkaline  K2MnO4  solution  with  a  few  drops  of  alcohol  (?). 
To  a  little  of  the  permanganate  solution,  acidified  with  H2SO4,  add  sul- 
phurous acid  (?).  Treat  the  acidified  solution  also  with  solutions  of  ferrous 
sulphate  and  oxalic  acid  (?). 

Reactions. 

(4)  i.  What  color  do  Mn-compounds  impart  to  a  borax  bead  in  the 
oxidizing  flame?  What  is  the  effect  of  the  reducing  flame?  2.  Heat  a 
little  of  an  Mn-compound  with  Na2CO3  and  KNO3  on  a  platinum  foil. 
What  does  the  resulting  mass  contain  ?  3.  To  a  little  of  the  solution  of 
the  chloride  in  water  add  (NH4)2S.  What  is  the  color  of  the  precipitate. 
Test  its  solubility  in  acids  (including  acetic  acid).  4.  Add  caustic  soda 
to  another  portion  of  the  chloride  solution.  Is  the  precipitate  soluble  in 
an  excess  of  the  reagent  ?  Is  its  color  affected  by  exposure  to  the  air  ? 
Explain. 

IRON.— Fe. 

(1)  Preparation. — Into   a  tube  of  Bohemian  glass  place  a  porcelain 
boat  filled  with  the  finely  powdered  oxide.     Pass  a  current  of  dry  H 
through   the  tube,  and  when  all  the  air  is  expelled  (how  could  you  test 
it?),  apply  heat  to  that  part  of  the  tube  which  contains  the  boat.     What 
is  formed  in  the  anterior  portion  of  the  tube  ?     After  a  red  heat  has  been 
maintained  for  10  minutes  allow  the  boat  to  cool  in  H,  and  examine  its 
contents.     Are  they  attracted  by  the  magnet  ?     Expose  the  product  to 
air  (?). 

How  is  iron  obtained  from  its  ores  on  a  large  scale  ?  What  are  its 
properties?  (see  Richter,  pp.  393  and  394).  Distinguish  between  cast- 
iron,  steel,  and  wrought-iron. 

Salts. 

(2)  Ferrous  Sulphate.— -Fe  SO4  -f  7H2O. 

To  25  grms.  of  Fe  in  the  form  of  nails  or  wire,  free  from  rust,  contained 
in  a  flask,  add  200  cc.  of  dilute  (i  :  4)  sulphuric  acid.  When  the  evolu- 
tion of  the  gas  (?  Note  its  odor  !)  is  no  longer  violent,  warm,  and  finally 
boil  until  the  liberation  of  gas  ceases.  A  sample  of  the  solution  poured 


IRON.  53 

into  a  test-tube  should,  on  cooling,  give  a  copious  separation  of  crystals. 
Filter  into  a  casserole  containing  2-3  cc.  of  cone,  H2SO4,  and  let  stand 
for  8  hours.  Collect  the  crystallized  product  in  a  funnel  the  stem  of 
which  is  closed  with  a  loose  plug  of  glass  wool,*  allow  the  mother 
liquor  to  drain  off,  wash  with  very  little  cold  water  (?),  and  dry  between 
sheets  of  filter  paper.  Examine  the  product  carefully.  Note  its  color, 
taste,  solubility  in  water  and  crystal  form.  What  other  salts  of  analo- 
gous composition  are  isomorphous  with  it  ? 

What  is  observed  when  some  of  the  salt  is  heated,  first  moderately,  then 
strongly,  in  a  tube  of  hard  glass  ? 

Expose  the  aqueous  solution  of  the  salt  to  the  air  for  several  hours  (?). 

(3)  Ferrous  Ammonium  Sulphate. — Fe  (NH4)2(SO4)2  -f-  6H2O. 

In  100  cc.  of  dilute  sulphuric  acid  dissolve  clean  iron  wire  till  no  more 
hydrogen  is  given  off;  neutralize  a  like  quantity  of  the  acid  exactly  with 
ammonia  water,  and  add  to  it  a  few  drops  of  dilute  sulphuric  acid.  Filter 
the  iron  solution  into  that  of  the  ammonium  salt.  Let  the  salt  crystallize, 
drain  it  on  a  funnel  provided  with  a  perforated  platinum  cone,  wash  and 
dry  as  described  under  (2).  Preserve  in  a  well-stoppered  bottle.  What 
metals  can  replace  the  iron  in  this  salt  without  altering  its  crystalline 
form? 

(4)  Ferric  Ammonium  Sulphate.— Fe2(SO4)3.(NH4)2SO4  -f  24H2O. 
Place  20  grms.  of  crystallized   ferrous  sulphate  into  a  porcelain   dish 

together  with  a  few  cc.  of  water  and  3.5  grms.  of  oil  of  vitriol.  Warm 
on  an  asbestos  plate,  adding  nitric  acid,  drop  by  drop,  until  no  further 
change  of  color  (?)  is  observed.  Evaporate  the  excess  of  HNO3,  dissolve 
the  residue  in  hot  water  and  add  3.5  grms.  of  (NH4)2SO4;  filter,  and  set 
the  solution  aside  for  crystallization.  Separate  the  crystals  from  the  mother 
liquor,  and  wash  and  dry  them  as  under  (2).  To  what  class  of  substances 
does  this  salt  belong  ?  Why  ? 

Reactions. 

(5)  In  a  borax  bead  dissolve  a  small  quantity  of  an  iron   compound, 
and   treat   it  successively  in   the   oxidizing  and  reducing  flames.     What 
changes  do  you  observe  ? 

(6)  Ferrous  Compounds. — Use  a  freshly  prepared  solution   of  ferrous 
sulphate  for  the  following  tests:    i.   To  a  few  drops  of  it,  diluted  with 
water,  add  ammonia.     Note  the  color  of  the  precipitate,  and  the  changes 
which  occur  on   exposure  to  the  air  (?).      2.   Add  (NH4)2S  to  another 
portion  (?).     Is  the  resulting  precipitate  soluble  in  HC1  ?    3.   In  a  porce- 
lain capsule  bring  together  a  little  of  the  ferrous  solution  and  a  drop  of 

*  It  is  better  to  use  a  perforated  platinum  cone,  and  to  remove  the  adhering  solution 
with  the  aid  of  a  filter  pump. 


54  EXPERIMENTS    IN    GENERAL   CHEMISTRY. 

a  potassium  ferrocyanide  solution.    Result  ?    4.  In  a  similar  manner  test  a 
drop  of  the  iron  solution  with  ferricyanide  of  potassium. 

(7)  Ferric  compounds.     In  the  presence  of  free  acids,  oxidizing  agents 
convert  iron  compounds  from  the  ferrous  into  the  ferric  condition,      i. 
Acidify  the  ferrous  sulphate  solution  with  sulphuric  acid,  warm,  and  add 
cone.  HNO3  until  it  fails  to  produce  a  change  in  color ;  the  iron  is  then 
in  the  ferric  state.      2.   Dilute  a  few  drops  of  the  yellow  liquid   with 
several  cc.  of  water  and  add  ammonia  (?).     3.  Test  a  drop  of  the  ferric 
solution  with  potassium  ferrocyanide  (?).     4.     Treat  a.  second  drop  with 
ferricyanide  of  potassium  (?).     5.   Mix  another  drop  with  a  solution  of 
potassium  sulphocyanate  (?).     6.   Conduct  H2S  into  some  of  the  ferric 
sulphate  solution.     What  do  you  observe?     Explain  the  reaction,  and 
write  the  equation  expressing  it.      7.   Place  a  piece  of  metallic  Zn  in  a 
test-tube  containing  a  solution  of  the  ferric  salt.     What  takes  place  ? 

(8)  Quantitative  estimation  of  iron.     Under  manganese  it  was  observed 
that  the  salt  potassium  permanganate  is  an  oxidizing  agent.     To  show 
how  this  salt  acts  with  iron  in  its  lower  form  of  oxidation,  fill  a  burette 
with  an  aqueous  solution  of  it ;  allow  it  to  drop  slowly  into  the  solution 
of  a  ferrous  salt  acidulated  with  H2SO4.     The  pink  color  of  the  perman- 
ganate immediately  disappears  on  stirring  with  a  glass  rod.     This  con- 
tinues until  the  ferrous  salt  is  completely  oxidized  to  the  ferric  state.     A 
drop  of  permanganate  added  in  excess  will  then  impart  a  faint  pink  color 
to    the   liquid.     This    indicates  that  the  reaction  is  ended.     Write  the 
equation. 

This  behavior  may  be  utilized  for  determining  the  quantity  of  iron  in  a 
solution.  That  this  may  be  done,  it  is  first  necessary  to  standardize  the 
FIG.  37-  permanganate  solution.  Proceed  as  follows:  Dissolve  about  2 
grms.  of  the  permanganate  in  1000  cc.  of  H2O.  Fill  a  burette 
'with  this  solution.  Weigh  out  .2  grm.  of  clean  piano  wire. 
Place  this  into  a  small  flask  (Fig.  37)  provided  with  a  cork 
and  valve.*  Cover  the  iron  wire  with  dilute  sulphuric  acid. 
Warm.  When  the  iron  is  completely  dissolved,  remove  the 
cork,  add  cold  water  to  the  solution,  and  slowly  admit  the  per- 
manganate until  the  final  pink  coloration  appears.  Note  the  volume  of 
the  K2Mn2O8  required  to  produce  this  effect.  Suppose  30  cc.  had  been 
consumed,  then  :  — 

30  cc.  K2Mn2O8  =  .2000    grm.  metallic  iron. 
I    "         «          =  .00666     "  " 

This  is  then  the  standard  of  the  permanganate  in  iron  units. 

*  With  a  sharp  knife  make  a  longitudinal  incision  of  about  I  cm.  length,  in  a  rubber 
tube,  and  close  one  end  by  means  of  a  glass  rod. 


COBALT   AND    NICKEL.  55 

Next,  dissolve  i  grm.  of  ferrous  ammonium  sulphate  in  100  cc.  distilled 
water,  add  5  cc.  H2SO4,  and  then  introduce  the  permanganate  until  the 
final  reaction  is  observed.  Calculate  the  percentage  of  iron  in  this  salt 
and  compare  the  experimental  result  with  the  theoretical  value. 

How  much  oxygen  will  each  molecule  of  K2Mn.2O8  give  up  in  oxidizing  ? 
How  many  molecules  of  FeO  can  be  changed  to  Fe2O3  by  a  molecule  of 
K2Mn208? 

COBALT.— Co  AND  NICKEL.— Ni. 
Reactions. 

i.  Dissolve  a  minute  quantity  of  a  cobalt  compound  in  a  borax  bead. 
Heat  first  in  the  oxidizing,  then  in  the  reducing  flame  (?).  2.  What  is  the 
behavior  of  nickel  compounds  under  like  conditions?  3.  Add  caustic 
alkali  to  a  solution  of  Co(NO3)2,  warm  the  mixture  (?).  What  action 
have  caustic  alkalies  on  solutions  of  nickel  salts?  5.  To  the  cobalt  solu- 
tion cautiously  add  ammonia.  After  a  precipitate  (?)  has  formed,  add 
more  of  the  reagent.  What  takes  place  ?  Expose  the  resulting  solution 
to  the  air  in  a  shallow  dish  (?).  6.  Treat  a  nickel  solution  in  an  analo- 
gous manner  (?).  7.  To  the  solutions  of  Co  and  Ni,  each  in  a  separate 
test-tube,  add  (NH4)2S.  Filter  and  wash  the  precipitated  sulphides,  and 
test  their  solubility  in  acids  (?). 

Note  the  colors  of  cobalt  and  nickel  salts,  in  the  hydrated  as  well  as  in 
the  anhydrous  state. 

Is  there  any  marked  difference  between  Co  and  Ni  in  respect  to  their 
chemical  deportment? 


Point  out  the  differences  in  the  reactions  of  Cr,  Mn,  Fe,  Co,  and  Ni. 

How  may  ferrous  compounds  be  distinguished  from  ferric  ?  What  con- 
ditions are  favorable  to  the  conversion  of  the  former  into  the  latter? 
The  latter  into  the  former  ? 

By  what  means  may  chromic  salts  be  changed  into  compounds  of 
chromic  acid?  How  may  the  reverse  change  be  effected? 

Devise  a  method  for  separating  the  elements  treated  in  this  chapter. 


Problems. — i.  How  much  K2Cr2OT  can  be  obtained  theoretically  from 
100  kilos  of  a  chromite  containing  58.6  per  cent,  of  Cr2O3?  2.  100 
grms.  of  a  pyrolusite  which  was  found  to  contain  4  per  cent,  of  impuri- 


56  EXPERIMENTS    IN    GENERAL   CHEMISTRY. 

ties,  will  give  what  volume  of  O,  measured  at  20°  C  and  745  mm.,  when 
strongly  ignited  ?  What  is  the  weight  of  the  residue,  assuming  that  one- 
half  of  the  impurities  was  moisture,  the  other  half  quartz  ?  3.  How 
many  grms.  of  HNO3  are  required  to  oxidize  12  grms.  of  crystallized  ferrous 
sulphate  ?  4.  What  percentage  of  metallic  iron  is  contained  in  a  salt,  of 
which  .7  grm.  are  exactly  oxidized  by  17.8  cc.  of  permanganate  solu- 
tion (standard  :  i  cc.  =  .0056  grm.  Fe)  ? 


APPENDIX. 


TABLE  OF  METRIC  WEIGHTS  AND  MEASURES, 


MEASURES  OF  LENGTH. 

I  metre  =  10  decimetres  —  100  centimetres  =  1000  millimetres, 
i  metre  —  1.09363  yards  =  3.2809  feet  =  39.3709  inches. 

MEASURES  OF  CAPACITY. 

i   cubic  metre  —  1000  litres  —  1,000,000  cubic  centimetres  —  1,000,000,000  cubic 
millimetres. 

i  litre  =  61.02705  cubic  inches  =  .035317  cubic  foot  3=  1.76077  pints  =  .22097  gallon. 

MEASURES  OF  WEIGHT. 
I  gram  =  weight  of  i  cc.  of  water  at  4°  C. 

I  Kilogram  =  1000  grams  =  100.000  centigrams  =  1,000,000  milligrams. 
I  Kilogram  =  2.20462  Ibs.  =  35-2739  ounces  =  15432.35  grains. 

TABLE  OF  ATOMIC  WEIGHTS  OF  ELEMENTS. 


Aluminium   . 

.  Al   .    . 

27.0 

Lead      .    .    . 

.  Pb  .    . 

....    207.0 

Antimony 

.  Sb  .    . 

I2O.O 

Magnesium 

Mg 

24.0 

Arsenic 

As 

7S.O 

Manganese 

Mn 

cc  o 

Barium      .    . 

.  Ba  .    . 

137.0 

Mercury 

•  Hg 

2OO.O 

Bismuth 

Bi 

.    .    .        208.0 

Molybdenum 

Mo      . 

Q6.O 

Boron    .    .    . 

.  B     .    . 

II.O 

Nickel    .    .    . 

.  Ni  .    . 

C.O.O 

Bromine    .    . 

.  Br  .    . 

.    .           .      80.0 

Nitrogen 

N 

14.0 

Cadmium 

Cd 

II2.O 

Oxygen 

o 

16.0 

Calcium    .    . 

.  Ca  .    . 

4O.O 

Phosphorus 

.  P    .    . 

31.0 

Carbon 

C 

I2.O 

Platinum 

Pt 

195.0 

Chlorine    .    . 

.  Cl    .    . 

35.5 

Potassium  .    . 

.  K   .    . 

....      39.0 

Chromium 

.  Cr   .    . 

C2.C, 

Silicon 

.  Si 

.  .         .     28.0 

Cobalt 

Co 

SO.O 

Silver 

Ag 

108.0 

Copper  . 

.  Cu  .    . 

63.3 

Sodium  .    .    . 

.  Na      . 

23.0 

Fluorine 

Fl 

IQ.O 

Strontium 

.  Sr      . 

87.5 

Gold  .    .    .    . 

.  Au      . 

.....   197.0 

Sulphur  .    .    . 

.  S    .    . 

....      32.0 

Hydrogen 

.  H       . 

i.o 

Tin         ... 

.  Sn       . 

118.0 

Iodine    .    .    . 

.  I     . 

.    .                127.0 

Zinc 

.  Zn     . 

....      65.0 

Iron 

Fe 

c6.o 

57 


UNIVERSITY  OF  CALIFORNIA  LIBRARY 
BERKELEY 


Return  to  desk  from  which  borrowed. 
This  book  is  DUE  on  the  last  date  stamped  below. 


DEC   3  19*' 


LD  21-100m-9,'47(A5702sl6)476 


C.  I 


